Blood And Buffers: Difference between revisions

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== Age ==
{| class="wikitable" style="color:black; background-color:#ddd; margin-left: auto; float:right"
| [[Biology]], [[Chemistry]]:
| Buffering Agents, Homeostasis
|-
| Grade Range:
| [[Middle School]], [[High School]]
|-
| Format:
| [[Hands-on]], [[Stage]]
|}


Elementary School, Middle School, High School
This demonstration on buffering agents will interest most students, especially when the connection is made to their own bodies!
 
== Format ==
 
Stage Show


== Materials ==
== Materials ==


    At least two 100 mL graduated cylinders or beakers
* Bromothymol Blue Indicator
    Bromthymol blue or universal indicator solution
* Water (DI Water ideal, but will work with Tap Water)
    Phosphate buffer solution, pH 7
* Dry Ice
    Dry ice – Keep in a cooler!
* Baking Soda
    Water – DI or tap water
* Phosphate Buffer Solution
    Pasteur pipettes and bulbs (or eye dropper / plastic transfer pipette)
* Three 500mL (Stage) or Four 100mL (Hands-on) Beakers
    Safety Goggles, Thermal Gloves, Nitrile Gloves, Lab Coats, Poster!
* Optional: Poster Board


== Safety Precautions ==
== Safety Precautions ==


Science Theatre demonstrators must keep the safety of themselves and their audience in mind at all times. All Science Theatre demonstrators must have read through the Safety Training page. The ST Safety Box with first aid kit, fire extinguisher, etc. should always be available to demonstrators. Always wear safety gloves, glasses, and a labcoat if handling chemicals; always perform potentially dangerous demonstrations at a safe distance from the audience; and always keep a very close eye on any volunteers you call from the audience. Demonstrators should wear goggles and lab coats, and the ST general safety box should be available to them. Dry ice is extremely cold! Please wear thermal gloves or AT LEAST nitrile gloves when placing pieces into beaker. Although the buffer solution is at pH 7, it still contains chemicals that we want to keep away from skin. Do not allow children to inhale any vapors or touch any solutions with their bare hands.
Please read the Dry Chemical, Liquid Chemical, and Cryogen sections of the [[Demonstration Safety]] page before performing this demonstration. The solutions for this demonstration can be poured down the drain, and do not need to go into a waste container.
 
[http://www.sciencelab.com/msds.php?msdsId=9927468: Bromothymol Blue MSDS]
 
[http://www.sciencelab.com/msds.php?msdsId=9925690: Phosphate Buffer Solution MSDS]
 
[http://education.jlab.org/frost/msds/dry_ice.pdf: Dry Ice MSDS]
 
== Preparation ==
 
Prepare enough pH 7 phosphate buffer solution as described below, plan on using 100 mL’s of buffer for each run through the experiment.
 
Purchase dry ice and keep in cooler
 
Prepare Phenolphthalein or universal indicator solution as described below. You will only need a few drops for each trial run, so if you make a lot, you can keep it and store it for the next acid/base demonstration
 
Try to acquire DI water – it contains less dissolved CO2 and therefore your control beaker will be closer to neutral pH. If not possible, tap water should be okay, just be aware its pH may range from 6 to 7.4, depending on the dissolved ions.
 
    PHOSPHATE BUFFER, pH 7
    Make 0.2 M NaH2PO4 solution
    137.9 g/mol *0.2 mol/L =  27.6 g/L
    Make 0.2 M Na2HPO4
    141.96 g/mol*0.2 mol/L = 28.4 g/L
    Mix together 500 mL of NaH2PO4 + 500 mL Na2HPO4 to make 1 L of phosphate buffer.
    Check the pH. Should be approximately 7. Bring up or down to 7 with NaOH or HCl, depending on the original pH.


BROMTHYMOL BLUE ALCOHOL SOLUTION
This demonstration requires: Cryogen Gloves, Goggles, Disposable Gloves
 
Add 0.5g of Bromthymol blue into 500ml of 95% ethanol and dissolve Add 500ml of distilled water Filter and store at room temperature


== Demonstration ==
== Demonstration ==


Put out two identical beakers, next to signs indicating BUFFERED or UNBUFFERED
;Preparation:
 
* Make the Phosphate Buffer Solution: In a 600mL beaker, add 13.8g of NaH<sub>2</sub>PO<sub>4</sub> to 500mL of DI water, and stir until completely dissolved. In a separate 600mL beaker, add 14.2g of Na<sub>2</sub> HPO<sub>4</sub> to 500mL of DI water, and stir until completely dissolved. Pour both solutions into a 1L bottle and mix. Label the solution as a 0.2M Phosphate Buffer Solution and write the date it was made. Be sure to test the pH of the solution to make sure it stays within two decimals of 7.0. If it is below, add a small amount of NaOH. If it is above, add a small amount of HCl.
To the unbuffered beaker, add 100 mL of DI water. To the buffered beaker, add 100 mL of the pH 7 phosphate buffer
* Setup the demonstration: In the first and last beakers, pour 100mL/500mL of water in, add 20 drops of the Bromothymol Blue indicator and stir. In the center beaker(s), add 100mL/500mL of the Phosphate Buffer Solution, then add 20 drops of the indicator and stir. All the solutions should be green, and should be the same shade as each other. Keep the buffer solution(s) in the center.
 
;Presentation:
Add 10 -20 drops of indicator solution – approximately 2 ml’s, the solutions should both be homogeneously green if using bromthymol blue, this indicates a pH somewhere between 6 and 7.6
# Show the Audience that the solutions look the same, and that the indicator you added shows that they are at or close to the same pH. Ask the audience what might happen to the first two solutions if you were to add a small amount of dry ice to them. What do they think it will do?
 
# Place a piece of dry ice into each solution, and watch what happens. The water solution will start turning yellow, while the buffer solution will stay unchanged! For the stage show, remove the dry ice from the buffer solution once the color change happens.
Using thermal gloves, break off a small chunk of dry ice and drop it into both beakers at the same time
# Ask the audience what they think might happen if you add baking soda to the last two solutions. What do they think it will do to the solutions?
 
# Add the baking soda to both solutions, and watch what happens. The water solution will start turning blue, while the buffer solution still stays unchanged!
Have the kids notice the difference in time it takes for a color change to occur in the two beakers. The buffered solution will change colors much more slowly than the unbuffered. If using bromthymol blue, acidic solution is indicated by a bright yellow color. If using universal indicator, once the solution acidifies the color may be yellow, orange or red, depending on how acidic the solution gets.
# For a hands-on show, use two beakers with the buffer solution so that one can have a piece of dry ice and the other can have the baking soda. You can leave all the solutions with their solvents; the buffer solutions should hold up to any changes for 30 minutes to an hour.
 
The chemicals used in the demonstration are safe to be poured down the drain of a sink – a hazardous waste container is not necessary.  
 
What to Say: Partner 1: Our cells are constantly undergoing chemistry. Cells must react oxygen with glucose from food to make energy to fuel all of your daily activities. (Ask what kind of things they do to need energy ie running, playing basketball, swimming….) In this process, carbon dioxide is released in the following reaction: (Point to the reaction on the poster)
 
C6H12O6 + 6O2 → 6CO2 + 6H2O +~38ATP
 
Partner 2: So every time you go for a run, your body breaks down the food you ate for breakfast, releasing carbon dioxide. The carbon dioxide, like we breathe out, can combine with water in the blood to form an acid – carbonic acid. (Point to this reaction on the poster) H2O + CO2 ⇔ H2CO3


Partner 1: Having all of this acid in our blood might start to hurt! Cells like to live in a very narrow pH range: from 6.8 to 7.4 – either just above or just below neutral pH (pH=7). (Depending on the age, you may need to state that pH is just a scale of how acidic or basic a solution is. Point to the pH scale on the poster). So how do we keep our cells happy and out of danger from all this acid when you go for a run?
== Why This Works ==


Partner 2: The key is buffers! Remember, a buffer is something that can resist a large change in pH when we add an acid or base. We have a bicarbonate buffer system in our blood, so when cells produce acid while working hard, the change in pH is absorbed rather than hurting our cells.
===Short Explanation===
''Buffer Solutions'' are able to resist a change in the pH, or acidity, which keeps the solution from becoming too acidic or basic. Our blood cells work to keep the blood in a stable pH range between 6.8 and 7.4, which is very narrow. One reason why our blood does this is because our bodies get energy from sugars and oxygen. When sugar reacts with oxygen, it breaks down into carbon dioxide (CO<sub>2</sub>) and water. CO<sub>2</sub> can react with water to make carbonic acid, which would make our blood more acidic, lowering the pH. We see this happen with the dry ice in the beaker with water, and how the indicator changes to yellow because of the acidity. Our blood prevents this by slowly reacting with the carbonic acid as it forms, and breaking it back down into CO<sub>2</sub> and water so that our lungs can breathe it out. Another reason why our blood has these buffering agents is because the carbonic acid can break down into bicarbonate. Bicarbonate can make the blood more basic, which increases the pH. We see this happen with the calcium carbonate in the beaker with water, and how the indicator turns blue because of it. Our blood prevents this by latching onto the extra bicarbonate so it doesn't react, and helping our kidneys filter it out.


Partner 1: Take, for example, these two cylinders. Both have water, an indicator (to tell us what the pH of the solution is), and are at approximately pH = 7. Now, I will add some dry ice, solid CO2. (Ask if anyone can hypothesize what will happen when we add the dry ice) Compare what happens to the beaker on the left versus the right. See how fast the color changes? That means that the pH is changing quickly – those cells are not happy! The other cylinder, however, is buffered, and any change in pH is being absorbed.
===Full Explanation===
''Buffer Solutions'' are able to resist changes in the pH, or acidity, which keeps the solution from becoming too acidic or basic. Our blood has a very narrow pH range that it needs to stay in, between a pH of 6.8 and 7.4, and to do so it needs to be able to buffer against any extra acid or base. If the blood becomes too acidic or basic, it will affect the bodies homeostasis, and can lead to illness or disease. The cells in our body get energy from simple sugars, with most of them using a reaction similar to this:


== Why It Is ==
;C<sub>6</sub>H<sub>12</sub>O<sub>6</sub> + 6 O<sub>2</sub> → 6 CO<sub>2</sub> + 6 H<sub>2</sub>O + ~38 ATP


The term pH is a simple way of expressing how acidic (or basic) a solution is. The standard pH scale runs from 0 to 14. If the pH is less than 7, the solution has more H+ than OH- and the solution is acidic. The lower the pH, the more H+ is present and the more acidic the solution is said to be. If the pH is greater than 7, the solution has more OH- than H+ and the solution is basic. The higher the pH, the more OH- is present and the more basic the solution is said to be. If the pH is equal to 7, the solution has the same amount of OH- as H+ and the solution is said to be neutral.
The extra CO<sub>2</sub> needs to be filtered out of our blood, and our bodies do that by using what is called the ''Carbonic Acid and Bicarbonate Buffer System''. The CO<sub>2</sub> is taken into the bloodstream, and some of it bonds with water to make carbonic acid. Carbonic acid is a weak acid, which means that most of it will dissociate into a bicarbonate and a free hydrogen, and some of it won't. Bicarbonate is the conjugate base to carbonic acid. This means that the amount of bicarbonate in the solution can change depending on how much carbonic acid is in the solution. The balance between the bicarbonate and carbonic acid is important for the pH of the blood, and works like this:


A buffer solution is a solution that will resist large pH changes upon the addition of an acid or base. Buffers resist pH change because there is an acid component to neutralize added base and a base component to neutralize added acid. The most common type of buffer in the chemistry laboratory is the combination of a weak acid like acetic acid and its conjugate base, an acetate salt. The conjugate acid-base pair differs by an H+ ion.
{| class="wikitable" style="color:black; background-color:#ddd; text-align: center; margin-right: auto"
|+ Carbonic Acid and Bicarbonate Buffer System
|Carbon Dioxide and Water
|
|Carbonic Acid
|
|Bicarbonate and Hydrogen
|-
|CO<sub>2 aq</sub> + H<sub>2</sub>O<sub> l</sub>
|'''⇋'''
|H<sub>2</sub>CO<sub>3 aq</sub>
|'''⇋'''
|HCO<sub>3</sub><sup>-</sup><sub>aq</sub> + H<sup>+</sup><sub>aq</sub>
|}


Indicators can be used to determine how acidic or basic a solution is. The one we are using today is a bromthymol blue, which turns blue in the presence of a base. As more acid is added to the solution, the indicator will change color to green at neutral pH and then is bright yellow at acidic pH’s, letting us know that the solution has changed its pH. Indicators change color due to the color of the different forms the chemical can take on. When it is protonated (say, in an acidic solution) the indicator turns yellow.
The red blood cells can move the extra CO<sub>2</sub> between the acid state, the base state and as dissolved CO<sub>2</sub> state, which keeps the system balanced and the pH neutral. This also allows the blood to keep the CO<sub>2</sub> in check until it can be breathed out by the lungs, or until it can be filtered out by the kidneys as bicarbonate or carbonic acid. When a buffer system isn't already in place against the changes, then the extra CO<sub>2</sub> can make the solution too acidic, as you saw with the dry ice in water. Likewise, if there isn't a way to take care of the extra bicarbonate, then the solution can become too basic, as you saw with the baking soda in water.


== Real Life Examples ==
== Additional Information ==


Alka seltzer tablets contain weak acids and their conjugate bases. When dissolved in water, you’ve made your own little buffer system! In biology labs, cells are often grown in buffers to keep them happy. Indicators are added to the solution to let the scientist know when he needs to change the media.
* Buffer solutions are a complicated topic, so do not be surprised if students are having trouble understanding the demonstration.
* This demonstration pairs well with the [[Blood Clotting]] demonstration.

Latest revision as of 15:51, 12 May 2016

Biology, Chemistry: Buffering Agents, Homeostasis
Grade Range: Middle School, High School
Format: Hands-on, Stage

This demonstration on buffering agents will interest most students, especially when the connection is made to their own bodies!

Materials

  • Bromothymol Blue Indicator
  • Water (DI Water ideal, but will work with Tap Water)
  • Dry Ice
  • Baking Soda
  • Phosphate Buffer Solution
  • Three 500mL (Stage) or Four 100mL (Hands-on) Beakers
  • Optional: Poster Board

Safety Precautions

Please read the Dry Chemical, Liquid Chemical, and Cryogen sections of the Demonstration Safety page before performing this demonstration. The solutions for this demonstration can be poured down the drain, and do not need to go into a waste container.

This demonstration requires: Cryogen Gloves, Goggles, Disposable Gloves

Demonstration

Preparation
  • Make the Phosphate Buffer Solution: In a 600mL beaker, add 13.8g of NaH2PO4 to 500mL of DI water, and stir until completely dissolved. In a separate 600mL beaker, add 14.2g of Na2 HPO4 to 500mL of DI water, and stir until completely dissolved. Pour both solutions into a 1L bottle and mix. Label the solution as a 0.2M Phosphate Buffer Solution and write the date it was made. Be sure to test the pH of the solution to make sure it stays within two decimals of 7.0. If it is below, add a small amount of NaOH. If it is above, add a small amount of HCl.
  • Setup the demonstration: In the first and last beakers, pour 100mL/500mL of water in, add 20 drops of the Bromothymol Blue indicator and stir. In the center beaker(s), add 100mL/500mL of the Phosphate Buffer Solution, then add 20 drops of the indicator and stir. All the solutions should be green, and should be the same shade as each other. Keep the buffer solution(s) in the center.
Presentation
  1. Show the Audience that the solutions look the same, and that the indicator you added shows that they are at or close to the same pH. Ask the audience what might happen to the first two solutions if you were to add a small amount of dry ice to them. What do they think it will do?
  2. Place a piece of dry ice into each solution, and watch what happens. The water solution will start turning yellow, while the buffer solution will stay unchanged! For the stage show, remove the dry ice from the buffer solution once the color change happens.
  3. Ask the audience what they think might happen if you add baking soda to the last two solutions. What do they think it will do to the solutions?
  4. Add the baking soda to both solutions, and watch what happens. The water solution will start turning blue, while the buffer solution still stays unchanged!
  5. For a hands-on show, use two beakers with the buffer solution so that one can have a piece of dry ice and the other can have the baking soda. You can leave all the solutions with their solvents; the buffer solutions should hold up to any changes for 30 minutes to an hour.

Why This Works

Short Explanation

Buffer Solutions are able to resist a change in the pH, or acidity, which keeps the solution from becoming too acidic or basic. Our blood cells work to keep the blood in a stable pH range between 6.8 and 7.4, which is very narrow. One reason why our blood does this is because our bodies get energy from sugars and oxygen. When sugar reacts with oxygen, it breaks down into carbon dioxide (CO2) and water. CO2 can react with water to make carbonic acid, which would make our blood more acidic, lowering the pH. We see this happen with the dry ice in the beaker with water, and how the indicator changes to yellow because of the acidity. Our blood prevents this by slowly reacting with the carbonic acid as it forms, and breaking it back down into CO2 and water so that our lungs can breathe it out. Another reason why our blood has these buffering agents is because the carbonic acid can break down into bicarbonate. Bicarbonate can make the blood more basic, which increases the pH. We see this happen with the calcium carbonate in the beaker with water, and how the indicator turns blue because of it. Our blood prevents this by latching onto the extra bicarbonate so it doesn't react, and helping our kidneys filter it out.

Full Explanation

Buffer Solutions are able to resist changes in the pH, or acidity, which keeps the solution from becoming too acidic or basic. Our blood has a very narrow pH range that it needs to stay in, between a pH of 6.8 and 7.4, and to do so it needs to be able to buffer against any extra acid or base. If the blood becomes too acidic or basic, it will affect the bodies homeostasis, and can lead to illness or disease. The cells in our body get energy from simple sugars, with most of them using a reaction similar to this:

C6H12O6 + 6 O2 → 6 CO2 + 6 H2O + ~38 ATP

The extra CO2 needs to be filtered out of our blood, and our bodies do that by using what is called the Carbonic Acid and Bicarbonate Buffer System. The CO2 is taken into the bloodstream, and some of it bonds with water to make carbonic acid. Carbonic acid is a weak acid, which means that most of it will dissociate into a bicarbonate and a free hydrogen, and some of it won't. Bicarbonate is the conjugate base to carbonic acid. This means that the amount of bicarbonate in the solution can change depending on how much carbonic acid is in the solution. The balance between the bicarbonate and carbonic acid is important for the pH of the blood, and works like this:

Carbonic Acid and Bicarbonate Buffer System
Carbon Dioxide and Water Carbonic Acid Bicarbonate and Hydrogen
CO2 aq + H2O l ⇋ H2CO3 aq ⇋ HCO3-aq + H+aq

The red blood cells can move the extra CO2 between the acid state, the base state and as dissolved CO2 state, which keeps the system balanced and the pH neutral. This also allows the blood to keep the CO2 in check until it can be breathed out by the lungs, or until it can be filtered out by the kidneys as bicarbonate or carbonic acid. When a buffer system isn't already in place against the changes, then the extra CO2 can make the solution too acidic, as you saw with the dry ice in water. Likewise, if there isn't a way to take care of the extra bicarbonate, then the solution can become too basic, as you saw with the baking soda in water.

Additional Information

  • Buffer solutions are a complicated topic, so do not be surprised if students are having trouble understanding the demonstration.
  • This demonstration pairs well with the Blood Clotting demonstration.
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