Think about all of the substances that are produced by your body. Did you know that these substances can cause imbalances in the body's pH? When body pH gets too high, alkalosis occurs. When pH gets loo low, we enter a state of acidosis. For example, a decrease in pH and an increase in CO2 will cause the body to have respiratory alkalosis which can be dangerous. Luckily, our bodies contain buffer solutions, and their role inside our body is to maintain pH at a stable value!

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Table of contents
    • This article is about buffers.
    • First, we will talk about acids, bases, and buffers.
    • Next, we will look at the importance of buffers in chemistry.
    • Then, we will dive into the different types of buffers that exist and give examples of buffers.
    • Lastly, we will take a look at calculations involving buffer solutions.

    Acids, Bases, and Buffers

    Now, you might be wondering what buffers are. A Buffer has the ability to resist changes in pH when an acid or a base is added to it.

    Buffers are solutions that maintain a constant pH when small amounts of acids or bases are added to them.

    To better understand the meaning of buffers, let's refresh our knowledge of acids and bases by looking at their definitions. Acids and bases have different definitions, created by different chemists and based on different properties.

    Arrhenius acids and basesProduce H+ ions in aqueous solutionsProduce OH- ions in aqueous solutions
    Bronsted-Lowry acids and basesProton (H+) donorsProton (H+) acceptors
    Lewis acids and basesElectron acceptorsElectron donors

    Arrhenius acids and bases can be divided by how they dissociate in water. Strong acids and bases dissociate completely in water, while weak acids and bases partially dissociate in water.

    When we talk about things in water it is almost always about Arrhenius acids, if you need acid in an organic solvent Bronsted comes into play. Lewis acids are very important for organic chemistry but are not part of the AP Exam material so you shouldn't worry about them for now.

    Importance of Buffers in Chemistry

    Buffers are very important in chemistry because many chemical reactions are pH-sensitive, meaning that they can only occur (favorably) under a narrow pH range.

    For example, most aquatic organisms can survive within a pH range of 6.5 - 8, and a pH outside of this range would lead to negative outcomes such as a decrease in reproduction, the appearance of diseases, and even death. Changes in pH can also affect the chemical state of pollutants such as copper and ammonia, causing an increase in toxicity to aquatic life.

    In our body, blood pH needs to be as close as possible to 7.4, and any changes in pH causing it to become higher or lower could lead to death. So, our blood plasma contains buffers to maintain a constant blood pH and keep it from becoming too acidic or basic.

    Generally, you would say that a Human can survive a blood pH of between 6.8 - 7.8 that is a narrow range indeed!

    Examples of Buffers

    • Carbonic acid, H2CO3 is a weak acid that dissociates in water and forms H+ ions and HCO3- ions. An example of a common buffer is an H2CO3/HCO3- buffer solution. This weak acid/conjugate base buffer system is very important to our bodies because it maintains our blood at a suitable pH of around 7.4. This is a roughly neutral buffer since it has a pH of over 7.
    • Another common buffer would be the sodium acetate/acetic acid buffer. This is often used by chemists in the lab since it is a slightly acidic buffer. It works between a pH of 3.6-5.
    • Finally, let us see a really basic buffer; A sodium carbonate/sodium bicarbonate buffer would work around a pH of 10.

    Types of Buffers

    First, let's talk about the composition of a buffer. Buffers are composed of either:

    • A weak acid + its conjugate base.

    • A weak base + its conjugate acid.

    A conjugate base is an acid that has lost a proton (H+).

    A conjugate acid is a base that has gained/accepted a proton (H+)

    A buffer solution can be either acidic or basic. A buffer's main job is to maintain the concentration of hydrogen (H+) and hydroxide (OH-) ions roughly the same, despite the addition of acidic or basic substances.

    When creating a buffer, the ideal ratio of weak acid/conjugate base or weak base/conjugate acid is 1 to 1. Having a one-to-one ratio allows the buffer the have a maximum buffer capacity, i.e the greatest resistance to pH changes.

    Acidic Buffers

    Acidic buffers are buffer solutions made up of a weak acid and its conjugate base.

    Let's use HF, a (super deadly) weak acid, as an example. If we wanted to make a buffer solution containing hydrofluoric acid, we would have to use HF and its conjugate base the fluoride anion. But, how would the fluoride be added?

    Normally, conjugate acids/bases are added to the solution as salts. In this case, a soluble salt such as sodium fluoride could be used since the sodium cation would not have an impact on the solution (a negligible cation). Therefore, by making a solution containing HF and sodium fluoride, we could create a HF/NaF buffer!

    Buffers prevent the formation of big amounts of H+ and OH- when acids or bases are added. So, if you add a strong acid or base to water containing a buffer, the acids won't react with the water - Instead, they will react with the buffer and not cause a great effect on the pH of the solution.

    Adding an Acid to an Acidic Buffer

    Usually, when an acid is added to water, it dissociates and forms H+ ions. This decreases the pH of a solution. But, when an acid is added to water containing an acidic buffer, the protons (H+ ions) from the acid will react with the anions (A-) formed from the ionization of the salt and form the weak acid HA.

    H+(aq) + A-(aq) HA (aq)

    Let's go back to our example involving an HF/F- buffer.

    When an acid is added to an aqueous solution containing HF/F- buffer, the fluorine anion (F-) from the buffer will react with the hydrogen ions from the acid and form HF.

    • F- reacts with the acid because F- is considered to be a better base than water and will therefore act as a proton acceptor instead of water.

    Adding a Base to an Acidic Buffer

    When a base is added to water, it forms OH- ions that lead to an increase in the pH of the solution. But, when a buffer is present in the solution, the hydroxide (OH-) ions will react with the weak acid in the buffer solution instead of reacting with water, forming water and A- ions.

    $$OH^{-}_{(aq)}+HA_{(aq)}\rightleftharpoons H_{2}O_{(l)}A^{-}_{(aq)}$$

    So, when a base is added to an aqueous solution containing an HF/F- acidic buffer, the HF will react with the OH- ions and form F- ions instead.

    • HF acts as the acid (H+ donor) instead of water because HF is a better acid than water!

    The pKa of Acidic Buffers

    When dealing with acidic buffers, you will need to know what their pKa or Ka is in order to be able to calculate the pH of a buffer solution. However, if you have a one-to-one ratio of weak acid/conjugate base buffer, then the pH of the buffer will be equal to its pKa.

    pKa is the negative log of Ka. In general, the higher the pKa, the weaker the acid will be.

    Ka is the acid dissociation constant.

    The pKa of a buffer also tells us the pH range of a buffer. For example, CH3COOH has a pKa of 4.8, so its buffer range would be between 3.8 and 5.8 (± 1 pH unit).

    The pH range of a buffer is referred to as the pH range in which a buffer can effectively function, neutralizing added acids and bases, and maintaining pH.

    Read "pH and pKa" to learn more about pKa and weak acids!

    Basic Buffers

    Basic buffers are buffer solutions made up of a weak base and its conjugate acid.

    How do you make a buffer solution containing a weak base? To make a buffer using a weak base, we need to mix it with its conjugate acid.

    For example, to make a buffer using NH3, we would need to add its conjugate base NH4+ by adding ammonium ion mixed with the negligible conjugate base of a strong acid. So, we could add NH4Cl, NH4Br, NH4I, and others to NH3 to make a buffer solution.

    If you need to recall the types of strong acids and bases, read "Acids and Bases".

    Adding an Acid to a Basic Buffer

    Let's pretend to have a weak base/conjugate acid buffer containing NH3 and its conjugate acid, NH4+. If we added acid to this buffer, what do you think would happen? If you guessed that the H+ ions from the acid would react with NH3 to form ammonium ion (NH4+), then you are on the right track!

    $$NH_{3\ (aq)}+H^{+}\rightleftharpoons NH_{4(aq)}^{+}$$

    By having the H+ reacting with ammonia instead of water, the buffer prevents the formation of H+ ions in water and therefore maintains a constant pH.

    • If a buffer was not present to maintain pH after the addition of an acid, pH would decrease due to an increase in the concentration of hydrogen ions in water.

    Adding a Base to a Basic Buffer

    On the other hand, If we added a base to the NH3/NH4+ buffer, the OH- ions from the base would react with the ammonium ion from the buffer to form NH3 and H2O.

    $$NH_NH_{4(aq)}^{+}+OH^{-}_{(aq)}\rightleftharpoons NH_{3\ (aq)}+H_{2}O_{(l)}$$

    Similarly, by having the OH- reacting with NH4+ instead of water, the buffer stops the formation of more OH- ions in water, preventing the pH from increasing/becoming more basic.

    It is important to realize that puffers always work in an equilibrium reaction. Why? Simply because they are better acids and bases but not that much better they can't be compared. If it was not an equilibrium reaction we would be talking about strong acids and bases.

    The Buffer Solutions Equation

    Before we dive into buffer calculations, let's review the basics of Ka. Ka is the equilibrium constant of the dissociation of a weak acid (HA).

    $$HA_{(aq)}\rightleftharpoons H^{+}_{(aq)}+A^{-}_{(aq)}$$

    Ka can be calculated by using the equilibrium expression. In this case it will look like this:


    Rearranging this expression to find H+ is pretty easy, looks like this:

    $$K_{a}\cdot \frac{[HA]}{[A^{-}]}=[H^{+}]$$

    Now, we can easily solve for hydrogen ion concentration, but it would be nicer to solve for pH wouldn't it? So let us take the negative logarithm of both sides. Keep in mind pH is just -log(H+) and pKa is -log(Ka)!

    $$pH=-log([H^{+}])=-log(Ka\cdot \frac{[HA]}{[A^{-}]})=-log(Ka)-log(\frac{[HA]}{[A^{-}]})=pKa+log\frac{[A^{-}]}{[HA]}$$

    Now this one has a name, it's called the Henderson-Hasselbalch equation, and is very handy for buffers as you will see in a second.

    Henderson-Hasselbalch Equation

    Now, let's look at an equation involving buffers. When dealing with buffers we can use the Henderson-Hasselbalch equation, which is given below. It's important to know that we can only use this equation when we have both a weak acid and its conjugate base in the solution.


    Since the Henderson-Hasselbalch equation is derived from Ka, you could also use the Ka expression to solve for pH. However, in buffers, using the Henderson-Hasselbalch formula is much more convenient!

    Let's look at an example!

    Find the pH of a buffer solution containing 0.15 M CH3COOH and 0.10 M CH3COONa.

    (Ka of acetic acid= 1.76 · 10-5)

    In this problem, acetic acid, CH3COOH is the weak acid, and the CH3COO- is the conjugate base (found in the sodium acetate salt, CH3COONa). Remember that conjugate bases are acids that lost an H+.

    Although the concentrations are not exactly the same, they are actually close to a 1 : 1 ratio, so this solution will function as a buffer.

    To calculate the pH of this buffer, we need to use the Henderson-Hasselbalch equation. But first, we need to convert Ka to pKa by using the formula: pKa= -log10 (Ka)

    $$pKa=-log_{10}(1.76\cdot 10^{-5})$$


    Now, we can plug all of the values into the Henderson-Hasselbalch equation to find pH.




    Now that you learned more about buffers, you should be able to tackle many other problems involving buffers!

    Buffers - Key takeaways

    • Buffers are solutions that maintain a constant pH when small amounts of acids or bases are added to them.
    • Buffers prevent the formation of big amounts of H+ and OH- when acids or bases are added by reaction with those acids and bases, instead of allowing them to react with water and cause changes in pH.
    • To calculate the pH of a buffer solution, we can use the Henderson-Hasselbalch equation.


    1. https://www.epa.gov/caddis-vol2/ph
    2. Malone, L. J., & Dolter, T. (2013). Basic concepts of of Chemistry. Hoboken, NJ: John Wiley.
    3. Salazar, E., Sulzer, C., Yap, S., Hana, N., Batul, K., Chen, A., . . . Pasho, M. (n.d.). Chad's general chemistry Master course. Retrieved May 4, 2022, from
    4. https://courses.chadsprep.com/courses/general-chemistry-1-and-2.
    Frequently Asked Questions about Buffers

    What are buffers and examples?

    Buffers are solutions that maintain a constant pH when small amounts of acids or bases are added to them.

    An example of a common buffer is an H2CO3/HCO3- buffer solution. This weak acid/conjugate base buffer system is very important to our bodies because it maintains our blood at a suitable pH of around 7.4.

    What are the 3 chemical buffers?

    Buffers can be acidic buffers and basic buffers. A third type simply does not exist.

    What is the function of the buffer?

    The function of a buffer solution is to maintain a constant pH when small amounts of acids or bases are added to them.

    Do buffers increase or decrease pH?

    Buffers do not increase nor decrease pH. A buffer has the ability to resist changes in pH (maintaining a constant pH).

    How do you tell if a solution is a good buffer?

    A good buffer is a solution that is able to resist changes in pH when an acid or a base is added to it.

    Test your knowledge with multiple choice flashcards

    Buffers are solutions that ________ when small amounts of acids or bases are added to them.

    _____ acids produce H+ ions in aqueous solutions

    Arrhenius bases produce ____ ions in aqueous solutions.


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