3 1.3 Acids and Bases

1.3.1 Introduction

Acids and bases serve many roles such as catalysts, reactants, or products in an organic reaction and as such, identifying and utilizing acids and bases correctly is critical in understanding organic chemistry reactions. This module focuses on the importance of acids and bases in organic chemistry reactions by focusing on the following skills: understanding the Lewis definition of acids, correctly identifying and differentiating between the types of acids and bases, identifying structural factors that determine relative strengths of acids and bases, using relative strengths of acids and bases to determine the position of equilibrium and to correctly predict the products of acid-base reactions.

Watch the following CrashCourse video for the key points of acidity related to organic chemistry.

1.3.2 Arrhenius Acids and Bases

Additional Resources: Arrhenius Acids and Bases

1. Arrhenius acids and bases by LibreTexts is available under a CC BY-NC-SA license.
2. Arrhenius acid-base reactions by LibreTexts is available under a CC BY-NC-SA license
3. Read and watch this video on Arrhenius acids and bases.

1.3.3 Bronsted-Lowry Acids and Bases

A Brønsted–Lowry acid is a substance that donates a hydrogen ion, H+, and a Brønsted–Lowry base is a substance that accepts a hydrogen ion. (The name proton is often used as a synonym for H+ because loss of the valence electron from a neutral hydrogen atom leaves only the hydrogen nucleus—a proton.) When gaseous hydrogen chloride dissolves in water, for example, a polar HCl molecule acts as an acid and donates a proton, while a water molecule acts as a base and accepts the proton, yielding chloride anion (Cl^–) and hydronium cation (H₃O⁺). This and other acid–base reactions are reversible, so we’ll write them with double, forward and backward arrows (scheme 1.3.1).

Scheme 1.3.1: Reaction Between An Acid And A Base

Chloride ion, the product that results when the acid HCl loses a proton, is called the conjugate base of the acid, and hydronium ion, the product that results when the base H₂O gains a proton, is called the conjugate acid of the base. Other common mineral acids such as H₂SO₄ and HNO₃ behave similarly, as do organic acids such as acetic acid, CH₃CO₂H.

Additional Resources: Bronsted-Lowry Acids and Bases

1. Read and watch this video on Bronsted Lowry acids and bases.
2. Brønsted-Lowry concept of acids and bases by LibreTexts is available under a CC BY-NC-SA license.
3. Bronsted-Lowry acids and bases may provide additional information in Introductory Chemistry, 1st Canadian Edition by David W. Ball and Jessie A. Kay is available under CC BY-NC-SA license.

Additional Resources: Conjugate Acid-Base Pairs

Read and watch this resource on conjugate acid-base pair.

1.3.4 Factors That Affect Acidity And Basicity

1. Watch the following videos that describe how acidity is related to structure.    Video 1    Video 2    Video 3     Video 4     Video 5
2. Watch this video that relates Ka and pKa to acidity.
3. Read this article that explain Ka and pKa; how this related to acid strength.
4. Read this page that lists a summary of factors that affect acid strength.

1.3.5 Predicting the Position of Equilibrium In Acid-Base Reactions

Now, that you have mastered the factors that affect acidity and basicity, you are ready to learn how to predict the position of the equilibrium in an acid-base reaction based on pKa values. H+ will always go from the stronger acid to the stronger base. That is, an acid will donate a proton to the conjugate base of a weaker acid, and the conjugate base of a weaker acid will remove a proton from a stronger acid.

Since water (pKa = 15.74) is a weaker acid than acetic acid (pKa = 4.76), for example, hydroxide ion holds a proton more tightly than acetate ion does. Hydroxide ion will therefore react to a large extent with acetic acid, CH₃CO₂H, to yield acetate ion and H₂O (scheme 1.3.2).

Scheme 1.3.2: Position of Equilibrium

Additional Resources:

1. Read about how pKa values can be used to predict the outcome of acid-base reactions.
2. This is another article that clearly describes the steps to determine the position of the equilibrium in acid-base reactions.

1.3.6 Lewis Acids and Bases

Although Bronsted theory is still widely used, it is restricted to compounds that can donate or accept protons. In 1923, G.N. Lewis from UC Berkeley proposed an alternate theory to describe acids and bases. His theory gave a more generalized explanation of acids and bases based on electron pairs. Through the use of the Lewis definition of acids and bases, chemists are now able to predict a wider variety of acid-base reactions. 

A Lewis acid is an electron pair acceptor while a Lewis base is an electron pair donor. A Lewis acid-base reaction occurs when a base donates a pair of electrons to an acid. A Lewis acid-base adduct, a compound that contains a coordinate covalent bond between the Lewis acid and the Lewis base, is formed. Lewis acids are also known as electrophiles (electron loving) and Lewis bases are known as nucleophiles (nucleus loving or positive charge loving). Lewis acids will have vacant orbitals to accept electrons. Lewis bases will have lone pairs or pi bonds to donate.  

There are different chemical species that can act as Lewis acids.

• All cations are Lewis acids since they are able to accept electrons. (e.g., Cu²⁺, Fe²⁺, Fe³⁺)
• An atom, ion, or molecule with an incomplete octet of electrons can act as an Lewis acid (e.g., BF₃, AlF₃). These are commonly used Lewis acids in several organic reations.
• Molecules where the central atom can have more than 8 valence shell electrons can be electron acceptors, and thus are classified as Lewis acids (e.g., SiBr₄, PF₅).
• Molecules that have multiple bonds between two atoms of different electronegativities (e.g., CO₂, SO₂)

Lewis bases are species that contain lone pairs such as NH₃ or H₂O or pi bonds such as CH₂=CH₂.

Some examples of acid-base reactions depicting the Lewis acid-base concept are shown here (scheme 1.3.3):

Scheme 1.3.3: Examples of Lewis Acid-Base Reactions

In reaction (a), HO^– ion is the nucleophile. Oxygen is the nucleophilic site and transfers a pair of electrons to carbon of the other molecule. Carbon carries a partial positive charge due to the polarity of the C-Br bond. Remember to cleave the C-Br bond or you will end up with more than an octet number of electrons in carbon’s valence shell.  The overall charge of both the reactants and the products remains the same, -1.

In reaction (b), the cyanide ion is the nucleophile and carbon is the nucleophilic site. It transfers a pair of electrons to the carbonyl carbon atom (this atom carries a partial positive charge due to the polarity of the C-O bond). In this case, the pi bond is broken and polarized towards the oxygen atom. The overall charge of the reactants and products is constant, -1.

In reaction (c), NH₃ is the nucleophile and nitrogen is the nucleophilic site. The electron pair on the nitrogen atom is donated to the empty p-orbital of the positively charged carbon atom. This carbon atom is the electrophilic site. The nitrogen atom in the product carries a formal charge of +1.

Additional Resources: Lewis Acids and Bases

1. A very detailed analysis of the Lewis acid-base theory can be found here.
2. Lewis acids and bases by LibreTexts is available under a CC BY-NC-SA license.

1.3.7 Drawing Curved Arrows

For a new organic chemistry student, drawing curved arrows can seem a complex process. However, the underlying principle is fairly simple, shows the movement of electrons from an electron rich site to an electron poor site. The following rules also help you draw the arrows correctly.

Other key points to note are:

1. The electronegativity of carbon atoms varies as Csp³ < Csp² < Csp. Do not move electrons from atoms of greater electronegtivity towards those of lower electronegativity.
2. Second row elements cannot have more than octet in their valence shells.
3. Watch this CrashCourse video on “Nucleophiles and Electrophiles”.