Introduction to Acids and Bases

Basic Introduction to Acids and Bases

Let's start our introduction to acids, bases, and salts by talking about the definition of acids and bases. There are different definitions associated with acids and bases, and each these definitions comes from different chemists.

Let's start with the Arrhenius definition of acids and bases.

Hydrochloric acid (with a formula HCl), for instance, is an Arrhenius acid because, in water, it separates into its individuals ions, increasing the H⁺ concentration in water.

$$\text{HCl }\stackrel{{\text{ H}}_2\text{O}}{\to }{\text{H}}^{+}{\text{ + Cl}}^{-}$$

On the other hand, sodium hydroxide (NaOH) is considered an Arrhenius base because, in water, it separates into OH^- and Na⁺, increasing the OH^- concentration in water.

$$\text{NaOH }\stackrel{{\text{ H}}_2\text{O}}{\to }{\text{Na}}^{+}{\text{ + OH}}^{-}$$

Next, we have the Brønsted-Lowry definition of acids and bases.

For example, in a chemical reaction between ammonia (NH₃) and hydrochloric acid (HCl), HCl will donate a H⁺ ion to NH₃ and become Cl^-, and NH₃ will accept a H⁺ ion to become NH₄⁺.

$$\underset{Base}{{\text{NH}}_3}\text{ + }\underset{Acid}{\text{HCl}}⟶{\text{NH}}_4^{+}{\text{ + Cl}}^{-}$$

Brønsted-Lowry acids are grouped bases of the amount of protons (H⁺) they are able to donate.

• Monoprotic acids can only donate one proton (H⁺). For example, HF, HCl and HClO₃.
• Diprotic acids can donate two protons. Diprotic acids include H₂SO₄ and H₂CO₃.
• Triprotic acids can donate three protons. A common example of a triprotic acid is H₃PO₄.

Introduction to Acids, Bases, and Salts

Now that we learned what acids and bases are, let's talk about salts! Basically, when an acid and a base react together, the products of the chemical reaction are salt and water. The chemical reaction between acids and bases is called a neutralization reaction.

The reaction below shows an example of a neutralization reaction between hydrochloric acid (HCl) and sodium hydroxide (NaOH), forming table salt (NaCl) and water (H₂O).

$$\underset{Base}{\text{NaOH}}\text{ + }\underset{Acid}{\text{HCl}}⟶\text{NaCl}{\text{ + H}}_2\text{O}$$

Some examples of common salts include baking soda (NaHCO₃), and bleaching powder (CaOCl₂).

Introduction to Acid and Base Concepts

When dealing with an introduction to acid and base, we need to explore some important concepts. One of them is is acid-base strength. The strength of acids and bases is based on how readily acids and bases perform their respective roles.

The strength of an acid depends on how well it dissociates and forms protons (H⁺ ions) in water. This means that a strong acid would more readily give up a proton than a weak acid, and a strong base more readily accepts a proton than a weak base.

• Strong acids will completely dissociate and yield nearly 100% H⁺ ions in water, while weak acids will partially dissociate in H₂O and not yield as many protons.

$${\text{HA}}_{(aq)}{\text{ + H}}_2{\text{O}}_{(l)}⟶{\text{H}}_{(aq)}^{+}{\text{ + A}}_{(aq)}^{-}$$

So, what happens to the strong acid after it has given up its proton? Because it wants to avoid accepting the proton back, it becomes a weak base! This is a trend that is seen in all strong acids and bases: they form a weak variant of the opposite once they have reacted in an acid-base reaction. We call these variants conjugates. Strong acids donate a proton and become a weak conjugate base. Weak acids donate a proton and become strong conjugate bases.

Similarly, the strength of a base depend on how well it dissociates in water, yielding hydroxide ions (OH^-). Strong bases accept a proton and become weak conjugate acids. Weak bases accept a proton and become strong conjugate acids.

Let's look at an example showing this! In chemical reaction below, HCl donates an H⁺ to the weak base (H₂O) and becomes a conjugate base, whereas the weak base (H₂O) accepts a proton from HCl and becomes a conjugate acid.

$$\underset{\text{Strong Acid}}{\text{HCl}}\text{ + }\underset{\text{Weak base}}{{\text{H}}_2\text{O}}⟶\underset{\text{Weak conjugate base}}{{\text{Cl}}^{-}}\text{ + }\underset{\text{Strong conjugate acid}}{{\text{H}}_3{\text{O}}^{+}}$$

For AP Chemistry, you should familiarize yourself with the following strong acids and bases (table 1).

Table 1. Strong acids and Strong bases.

The second concept we need to kind in mind is pH and the pH scale.

Figure 1 shows the pH scale. The pH scale is used by chemists to measure the tendency of molecules to be acidic or basic. On the left, we have the quality of "acidic," and on the right, we have "basic". The pH scale scale ranges from 0 up to 14, with zero being extremely acidic, 14 being extremely basic, and 7 representing a neutral solution.

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For example, water (H₂O) has a pH of 7, so we know it’s neutral. Black coffee, however, has a pH of 5, meaning that it is a weak acid. Baking soda, with a pH of around 9.5, is considered a weak base!

When strong base is added to a strong acid, the recently added OH^- ions will react with the dissociated H⁺ ions from the strong acid, changing the concentration of H⁺ ions in solution and causing the pH to change!

However, if you added a strong base to a weak acid, the pH would not undergo major changes because the strong base would mostly react with the HA molecules that did not dissociate with its ions.

Introduction to Titration of Acids and Bases

Now, let's make an introduction to titration of acids and bases, and explore the basic concepts related to it.

This means that if we wanted to figure out an acid's unknown concentration, we would use a base with a known concentration, and slowly add it to the acid until the mixture is neutralized. If we know the exact amount of base that we added, as well as its concentration, we can deduce the acid's unknown concentration!

Figure 2 shows the common laboratory setup used for acid/base titrations. The solution of unknown concentration is usually placed in a flask and a few drops of an indicator is added to it. The solution of known concentration is placed in the buret and added drop-wise to the sample in the flask until the solution changes color.

Fluid Electrolyte and Acid Base Balance Introduction to Body Fluids

To finish off, let's explore fluid, electrolyte and the acid-base balance of homeostasis. Although this is more important in biology and you most likely won't encounter it in your chemistry exam, it is a very important topic!

In the body of an average adult, there is an average of 40 L of body fluids (figure 3). The intracellular fluid is the fluid inside body cells, and it consists of mostly water and electrolytes like potassium (K⁺), magnesium (Mg²⁺), and HPO₄^2-. The extracellular fluid is found outside body cells and contains electrolytes such as Na⁺, Cl^-, HCO₃^- and Ca²⁺.

One of the many function of electrolytes is to help maintain acid-base balance in homeostasis. In this case, acids are considered electrolytes that release H⁺ ions in water, while bases are electrolytes that release OH^- ions in water.

If you understand the concepts discussed in this explanation, you'll have a really strong foundation that will help you during the AP chemistry test and also in higher-level chemistry courses!