Buffers

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!

• 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.

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 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.

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.

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 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.

$H^{+}\left(aq\right)+A^{-}\left(aq\right)\underset{}{\underset{}{\rightleftharpoons }}HA\left(aq\right)$

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.

When dealing with acidic buffers, you will need to know what their pKa or K_a 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 pK_a.

The pK_a 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).

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 NH_3, 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.

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.

The Buffer Solutions Equation

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

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

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

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

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)!

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

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.

$$pH=pKa+log\frac{[A^{-}]}{[HA]}$$

Let's look at an example!

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