- Acid/base questions
- Acid-base definitions
- Chemistry of buffers and buffers in our blood
- Ka and acid strength
- Autoionization of water
- Definition of pH
- Strong acid solutions
- Strong base solutions
- Weak acid equilibrium
- Weak base equilibrium
- Relationship between Ka and Kb
- Acid–base properties of salts
- pH of salt solutions
- Common ion effect and buffers
- Buffer solutions
- Buffer solution pH calculations
Chemistry of buffers and buffers in our blood
Let’s go to a chemistry lab and conduct a simple experiment. As shown below, we have
- two solutions - 50 mL of A and 50 mL of B respectively
- a solution of 0.2M hydrochloric acid (HCl)
- a solution of 0.2M sodium hydroxide (NaOH)
- pH meter to measure pH of the solution
Illustration of experiment 1 and experiment 2
Experiment 1: The pH of solution A is 7.0 i.e. it’s neutral. When we add 10 mL of 0.2M HCl to it, the pH decreases to 1.5. On the other hand, when we add 10 mL of 0.2M NaOH to solution A the pH shoots up to 12.5. In both these cases we see a drastic change in pH due to either the increase or decrease of proton [H] concentration. Remember, pH of a solution is dependent on the concentration of hydronium ions or simply put as [H].
Experiment 2: Let’s see what happens when we repeat the same experiment with solution B, whose pH is also maintained at 7.0. When we add 10 mL of 0.2M HCl to it, the pH decreases by only 0.2 units to 6.8. Next, when we add 10 mL of NaOH to solution B, the pH just slightly rises to 7.2 from 7.0. In both these cases we do not see a drastic change in pH, as we observed with solution A.
So, our observation from this experiment is that solution A underwent drastic changes in pH upon addition of a strong acid or a strong base, while solution B resisted a change in pH.
Can you guess what the difference might be between solutions A and B?
Solution A is pure ‘water’, while solution B is a ‘buffer’.
How do we define a buffer?
“A buffer is an aqueous solution that resists changes in pH upon the addition of an acid or a base”. Also, adding water to a buffer or allowing water to evaporate from the buffer does not change the pH of a buffer significantly. Buffers basically constituent a pair of a weak acid and its conjugate base, or a pair of a weak base and its conjugate acid (as will be discussed next).
How do we prepare a buffer?
A buffer is essentially prepared in two ways
- mixing a large volume of a weak acid with its conjugate base (eg. acetic acid – acetate ion, CHCOOH – CHCOO)
- mixing a large volume of weak base with its conjugate acid (eg. ammonia – ammonium ion, NH – NH)
Phosphate buffer is a very commonly used buffer in research labs. It falls under the second category. It's made up of a weak base (HPO) and its conjugate acid (HPO). The pH of a phosphate buffer is usually maintained at a physiological pH of 7.4.
How does a buffer work?
Buffer, as we have defined, is a mixture of a conjugate acid-base pair that can resist changes in pH when small volumes of strong acids or bases are added.
- When a strong base is added, the acid present in the buffer neutralizes the hydroxide ions (OH).
- When a strong acid is added, the base present in the buffer neutralizes the hydronium ions (HO).
Let’s understand this principle through the two examples we listed above.
- Suppose we have a buffer that contains acetic acid (CHCOOH) and its conjugate base, acetate ion (CHCOO), as depicted below.Illustration of beaker containing acetic acid and its conjugate base, acetate ionWhen a strong acid is added, the acetate ions neutralize the hydronium ions producing acetic acid (which is already a component of the buffer).Chemical reaction diagram of acetate ions neutralizing hydronium ions producing acetic acidOn the other hand, when a strong base is added, acetic acid consumes the hydroxide ions producing acetate ions (which is also already a constituent of the buffer).Chemical reaction diagram of acetic acid consuming hydroxide ions producing acetate ions
- Suppose we have a buffer that contains a weak base (NH) and its conjugate acid, ammonium ion (NH), as shown below.Illustration of beaker containing weak base NH3 and its conjugate acid, ammonium ionWhen a strong acid is added, ammonia neutralizes the hydronium ions producing ammonium ions (which is already a component of the buffer).Chemical reaction diagram of ammonia neutralizing hydronium ions producing ammonium ionsOn the other hand, when a strong base is added, ammonium ions consume the hydroxide ions producing ammonia (which is also already a constituent of the buffer).Chemical reaction diagram of ammonium ions consuming hydroxide ions producing ammonia
So now we have an understanding of how a buffer neutralizes the hydronium ions or the hydroxide ions to resist a change in the pH upon the addition of a strong acid or a strong base respectively.
Buffer capacity of a buffer
Let’s assume our buffer is made up of a weak acid (HA) and its conjugate base (A). In this case, the equilibrium constant of the weak acid will be represented as:
The above expression can be rearranged to give:
Since is a constant, [HO] will depend directly on the ratio of [HA]/[A].
The function of a buffer is to keep the pH of a solution within a narrow range. As you can notice from the above equation, the ratio of [HA]/[A] directly influences the pH of a solution. In other words, the actual concentrations of A- and HA influence the effectiveness of a buffer.
The more A and HA molecules available, the less of an effect the addition of a strong acid or base will have on the pH of the solution. For example, let’s see what will happen if we add a strong acid such as HCl to this buffer. Initially, the protons produced will be taken up by the conjugate base (A)
This will slightly change the pH by altering the ratio [HA]/[A] as [A] and [HA] are constantly changing, but as long as there is enough A present, the change in pH will be small. But if we keep adding HCl, eventually A will run out. Once there is no more A left, any additional HCl will donate its proton to water (HCl + HO → HO + Cl). This will dramatically increase the concentration [HO], leading to a drastic change in the pH of the solution.
So, in order to be an effective buffer,
- The number of moles of the weak acid and its conjugate base must be significantly large compared to the number of moles of strong acid or base that may be added.
- The best buffering will occur when the ratio of [HA] to [A] is almost 1:1. In that case pH = pK. Buffers are considered to be effective when the ratio of [HA] to [A] ranges anywhere between 10:1 and 1:10.
Buffering system of blood
Maintaining a constant blood pH is critical for the proper functioning of our body. The buffer that maintains the pH of human blood involves a carbonic acid (HCO) - bicarbonate ion (HCO) system.
Diagram carbonic acid - bicarbonate ion system in human blood
When any acidic substance enters the bloodstream, the bicarbonate ions neutralize the hydronium ions forming carbonic acid and water. Carbonic acid is already a component of the buffering system of blood. Thus hydronium ions are removed, preventing the pH of blood from becoming acidic.
Chemical reaction diagram of bicarbonate ions neutralizing hydronium ions forming carbonic acid and water
On the other hand, when a basic substance enters the bloodstream, carbonic acid reacts with the hydroxide ions producing bicarbonate ions and water. Bicarbonate ions are already a component of the buffer. In this manner, the hydroxide ions are removed from blood, preventing the pH of blood from becoming basic.
Chemical diagram of carbonic acid reacting with hydroxide ions producing bicarbonate ions and water
As depicted below, in the process of neutralizing hydronium ions or hydroxide ions, the relative concentrations of carbonic acid (HCO) and bicarbonate ions (HCO) fluctuate in the blood stream. But this slight change in the concentrations of the two components of the buffering system doesn’t have any adverse effect; the critical thing is that this buffering mechanism prevents the blood from becoming acidic or basic, which can be detrimental.
Diagram of blood pH maintained at approx. 7.4 by the carbonic acid – bicarbonate ion buffering system
The pH of blood is maintained at ~ 7.4 by the carbonic acid – bicarbonate ion buffering system.
Why is it so critical to maintain the pH of our blood?
Believe it or not, if our blood pH goes to anything below 6.8 or above 7.8, cells of the body can stop functioning and the person can die. This is how important the optimum pH of blood is!
Enzymes are very specific in nature, and function optimally at the right temperature and the right pH; if the pH of blood goes out of range, the enzymes stop working and sometimes enzymes can even get permanently denatured, thus disabling their catalytic activity. This in turn affects a lot of biological processes in the human body, leading to various diseases.
Want to join the conversation?
- So, as I was reading the last few sentences about detrimental consequences of a change in pH of our blood. I thought, could this be related to why patients who have ketoacidosis feel short of breath or tachypneic? So I did a bit of reading and apparently the increased breathing rate is the body's attempt to breath out as much CO2 as possible. I thought I would share because it is a interesting example of the importance of the blood buffering system and how it is linked to respiration.(147 votes)
- Yes, great connection. Excess C02 in the blood results in acidosis or low blood pH. In order to compensate of this low pH the body must get rid of the excess C02 and it does this by hyperventilating or high breathing rate. Keep it up!!(65 votes)
- A change in blood pH causes which diseases?(8 votes)
- Basically, (the pun a happy coincidence) alterations in blood ph cause acidosis or alkalosis, which is simply saying the same thing as alterations in blood chemistry. These arent disease states themselves, but may be the results of other disease states, or ingesting of certain substances. The alterations in blood chemistry can cause fatal complications, by affecting the oxygen disassociation curve, breathing (if respiratory is not the initial source of the imbalance) and the equilibrium of countless reactions needed to maintain body homeostasis(30 votes)
- Good day,
I would like to know how this carbonic acid-bicarbonate buffer system could be applied to the regulation of blood pH following acidic drug consumption.(5 votes)
- Under the section Buffer capacity of a buffer. I am very confused about the equation being used for Ka. Where is the H3O+ coming from, and how do you know where to plug in the acid/base? Can you please explain this in greater detail. Thanks.(1 vote)
- Ka is the dissociation constant for an acid, where we have products over reactants.
HA + H2O --> A- + H3O+
acid + water --> conjugate base + hydronium ion
Ka = [A-][H3O+]/[HA]
**Note that pure liquids and solids are excluded from all K constants.(5 votes)
- Why have a weak acid buffer a solution as opposed to a weak base?(3 votes)
- My knowledge about chemistry or biochemistry is extremely poor. Is there any way I can pass the biochemistry class?(2 votes)
- What if in a buffer solution, we increase volume of HCl or concentration HCl? does it both increase the pH of the solution dramatically when it exceeds the buffering capacity?
I'm little confuse because I remember that volume would decrease the concentration because it is mol/L. so wouldn't increase the volume of hcl to buffer solution would means decrease the concentration of HCl so the buffer solution would resist the change in pH?(2 votes)
- Is there a way to make these note pages printable? I'd like to print them for study use(1 vote)
- As far as I know, there isn't except manually zooming out and hitting the "print screen" button and then pasting it to "paint", etc..(1 vote)
- what is the chemistry of acid-base buffer in human blood(1 vote)
- How do phosphate and protein buffers work?(1 vote)