Properties of Buffers

Our stomachs contain gastric acid to break down food. This acid can even dissolve steel! How is it that such a potent acid is in our stomachs? Well, our body has a "safety" in place, so our stomach isn't acidic enough to dissolve itself! Stomach cells produce bicarbonate which buffers the acid to prevent it from getting too strong (which we will discuss below).

The buffer keeps the stomach pH stable (which we will also discuss later). This pH is between 1.5-3.5. It keeps stomach acid from getting too acidic, but it also prevents it from getting so basic that it can no longer dissolve food. In this article, we'll be diving into buffers and see about their preparation, their properties, and their applications.

• This article is about the properties of buffers.
• First, we will cover buffer properties.
• Then we'll use the Henderson-Hasselbalch equation to calculate the pH of a buffer solution.
• Thereafter, we'll learn about what is happening in a buffer solution.
• Next, we'll walk through buffer preparation.
• Lastly, we'll cover what makes a good buffer and buffer applications.

pH Measurements and the Properties of Buffers

I talked a lot about buffers and pH, but what are they?

Now that we know the basic definitions, let's cover the properties of buffers.

Properties of Buffers

Buffers have an identifying set of characteristics, these are:

1. A definite pH

2. pH won't change over time

3. Dilution won't change pH

4. A slight pH change if a small amount of acid/base is added

5. A range where they are the most effective

As mentioned previously, buffers have a buffer capacity, which is the amount of acid/base that can be added until a significant pH change occurs. Typically, we consider this to be a change of 1 unit. (ex. from 3 to 4). A buffer's concentration determines its capacity. As concentration increases, more acid/base can be absorbed until a large change happens.

Figure 1: Buffer Capacity - the greater the concentration, the greater the buffer capacity. Study Smarter Orginal

As illustrated above, the greater the concentration of the buffer, the greater the capacity. The "spike" in the graph shows where the pH begins to significantly change (buffer isn't as effective). After the spike, it becomes an effective basic buffer.

The buffer range is the range of pH values where a buffer is most effective. The best solutions have a 50:50 ratio of acid to base. Generally, a buffer isn't effective when one component is >10% of the other. For example, a buffer is prepared that has 0.1 M acetic acid (CH₃COOH) and 0.008 M acetate (CH₃COONa). If more base is added, the buffer will be "overwhelmed" and the pH will change significantly.

Figure 2: Once there is a 10:1 ratio of base to acid, the pH begins to change significantly as more base is added. StudySmarter Original.

As more base is added, the acetic acid converts to acetate. Once there is a 10:1 ratio of acetate to acetic acid (shown by the black dot), the pH starts changing significantly. From the data, we see that the acetic acid buffer loses effectiveness around pH = 6.

pH Measurements and the Henderson-Hasselbalch Equation

We measure a buffer's pH using the Henderson-Hasselbalch equation.

We measure the pH to see if our buffer is "working". If we add an acid/base, we want to ensure a minimal pH change, or else we would switch buffers.Let's say we have a buffer solution of CH₃COOH (weak acid) and CH₃COONa (conjugate base). Our acid equilibrium looks like:$$CH_3COOH\leftrightarrow CH_3COO^-+H^+$$

When a strong acid is added, its H⁺ ions combine with the CH₃COO^- ions and re-form the weak acid. This "neutralizes" the acid, preventing a large pH change.

Structure and Properties of Water pH Buffers

There are two types of buffers: weak acids + their conjugate base, or weak bases + their conjugate acid. Buffers are aqueous solutions (dissolved in water).

The previous example was a weak acid and its conjugate base. (acetic acid/acetate). We covered what happens when this type has a strong acid added to it, but what about a strong base?

\(CH_3COOH\leftrightarrow CH_3COO^-+H^+\)

\(CH_3COOH+OH^-\leftrightarrow CH_3COO^-+H_2O\)

The base adds OH^- ions which react with the weak acid to produce water, so the pH is unchanged.

Now for weak base/conjugate acid pairs. The main difference between the types is there is a weak base equilibrium, instead of a weak acid one. Let's look at the NH₄OH/NH₄Cl pair as an example. The equilibrium equation is:

$$NH_4OH\leftrightarrow NH_4^+ + OH^-$$

If an acid is added, the H⁺ ions combine with the weak base to form water

$$NH_4OH+H^+\leftrightarrow NH_4^+ +H_2O$$

When a base is added, the OH^- ions will react with the NH₄⁺ ions to form the weak base.

$$NH_4^+ + OH^-\leftrightarrow NH_4OH^+$$

Now that we've covered how a buffer works, we'll cover how to make and choose a buffer.

Determining an Appropriate Buffer

When choosing a buffer, you want to ensure that it can buffer the system you want. The buffer range should cover the pH you wish to stabilize. Some common buffers are:

For example, if you want to keep a system neutral (pH=7), you would choose the phosphate buffer (phosphate + phosphoric acid).

Preparing a Buffer Solution

Once you have chosen your buffer, you can prepare the solution. Let's walk through the steps:

Step 1: Use the Henderson-Hasselbalch equation to calculate the acid/base concentration.

Plug in your desired pH and the buffer pK_a into the equation, solving for the ratio of base to acid.

Step 2: Mix the required amounts of the weak acid/base with its conjugate base/acid.

The conjugate base can be added directly to the acid or a strong base can be added to the weak acid, partially converting it to the conjugate base (the same can be done with a strong acid+weak base)

Step 3 (optional): Adjust the pH.

When experimenting, there may be some error that makes the pH slightly off. Using a pH meter, you can monitor the pH as you add small amounts of strong acid/base to get the desired pH.

Characteristics of a Good Buffer

What makes a buffer "good"?

• [HA] = [A^-]. The logarithm in the Henderson-Hasselbalch equation will be equal to zero, so the pH = pK_a. This means the system can equally absorb acids/bases.
• If the desired pH is basic, then a weak base + its acid is the most effective. It's the opposite for acids.
• The desired pH is within the buffer's range

Applications of a Buffer Solution

Buffers have several applications. Enzymes must be kept in a specific pH range or else they will degrade\lose usefulness, so the body uses buffers to stabilize the pH. Buffers are also important in manufacturing and industry. They are used during fermentation and during the production of products like inks, dyes, and paper. They are also important for food preservation.