Premed · Premed · Organic Chemistry 1
Lecture 2: Acids and Bases in Organic Chemistry
Organic Chemistry I
Learning Objectives
By the end of this lecture, students will be able to:
- Define acids and bases according to Bronsted-Lowry and Lewis theories
- Identify conjugate acid-base pairs in organic reactions
- Use pKa values to predict the direction of acid-base equilibria
- Explain the structural factors that influence acidity and basicity
- Predict relative acid strength based on element effects, induction, resonance, and hybridization
- Recognize acid-base reactions as a foundation for understanding organic reaction mechanisms
Lecture Content
I. Bronsted-Lowry Acid-Base Theory
A Bronsted-Lowry acid is defined as a proton (H+) donor, while a Bronsted-Lowry base is a proton acceptor. In every Bronsted acid-base reaction, a proton is transferred from the acid to the base, generating a conjugate acid-base pair on each side of the equation. When an acid donates its proton, the species that remains is the conjugate base; when a base accepts a proton, the resulting species is the conjugate acid. For example, in the reaction of acetic acid with water (CH3COOH + H2O), acetic acid loses a proton to become the acetate ion (its conjugate base), while water gains a proton to become the hydronium ion (its conjugate acid).
Some substances are amphoteric, meaning they can function as either an acid or a base depending on the reaction partner. Water and alcohols are common examples. Proton transfer reactions are pervasive throughout organic chemistry, appearing in the protonation of leaving groups, the deprotonation of carbon acids to form carbanions or enolates, and the acid-base catalysis that drives many biological processes.
II. The Equilibrium Constant and pKa
The strength of an acid is quantified by its acid dissociation constant, Ka. For the dissociation HA giving H+ and A-, Ka equals the product of the concentrations of H+ and A- divided by the concentration of HA. A larger Ka indicates a stronger acid. Because Ka values span many orders of magnitude, chemists use the more convenient pKa scale, defined as the negative logarithm of Ka. A lower pKa corresponds to a stronger acid, while a higher pKa indicates a weaker acid whose conjugate base is correspondingly stronger.
The pKa scale in organic chemistry covers an enormous range. Strong acids such as HI (pKa approximately -10) and sulfuric acid (pKa approximately -3) sit at one extreme. Carboxylic acids fall around pKa 4-5, water and alcohols around 15-16, ketone alpha-hydrogens near 20, terminal alkyne C-H bonds near 25, alkene C-H bonds near 44, and alkane C-H bonds near 50 at the other extreme. To predict the direction of an acid-base equilibrium, one compares the pKa values of the acids on each side: equilibrium favors formation of the weaker acid, which is the species with the higher pKa. The equilibrium constant can be calculated as Keq = 10 raised to the power of (pKa of the product acid minus pKa of the reactant acid).
<image>A pKa scale displayed as a vertical number line ranging from -10 to 50, with representative organic acids placed at their approximate pKa values. Includes: HI (-10), H2SO4 (-3), H3O+ (-1.7), carboxylic acids (~5), ammonium ion (~9-10), phenol (~10), water/alcohol (~15-16), alpha-hydrogen of ketones (~20), terminal alkynes (~25), hydrogen gas (~35), ammonia (~38), alkenes (~44), and alkanes (~50). An arrow on the left indicates "increasing acid strength" going downward and "increasing conjugate base strength" going upward.</image>
III. Lewis Acid-Base Theory
The Lewis theory offers a broader framework for understanding acid-base chemistry. A Lewis acid is any species that accepts an electron pair and therefore must possess an empty or partially empty orbital. Examples include BF3, AlCl3, the proton (H+), carbocations, and metal cations such as Zn2+ and Mg2+. A Lewis base is any species that donates an electron pair, typically through a lone pair or pi electrons. Common Lewis bases include ammonia, water, alcohols, ethers, halide ions, and alkenes.
Every Bronsted-Lowry acid-base reaction is also a Lewis acid-base reaction, but the reverse is not true. The Lewis framework explains reactions that involve no proton transfer at all, such as the coordination of boron trifluoride with diethyl ether to form an adduct, the binding of metal catalysts to organic substrates, and the attack of a nucleophile on an electrophilic carbon. The curved arrow notation used throughout organic chemistry directly reflects Lewis acid-base interactions: the arrow always points from the electron-rich species (the Lewis base) toward the electron-poor species (the Lewis acid).
IV. Factors Affecting Acidity: Element Effects
Two periodic trends govern how the identity of the atom bearing the acidic hydrogen influences acid strength. Across a row of the periodic table (left to right), acidity increases as electronegativity increases. Thus CH4 (pKa approximately 50) is far less acidic than NH3 (pKa approximately 38), which is less acidic than H2O (pKa approximately 15.7), which is less acidic than HF (pKa approximately 3.2). More electronegative atoms better stabilize the negative charge on the conjugate base, making the acid stronger.
Down a column of the periodic table, acidity also increases, but for a different reason. HF (pKa approximately 3.2) is much less acidic than HCl (pKa approximately -7), which is less acidic than HBr (pKa approximately -9), which is less acidic than HI (pKa approximately -10). As atom size increases, bond strength to hydrogen decreases, and the larger atom can spread the negative charge over a greater volume through its higher polarizability. This size and polarizability effect dominates over electronegativity when comparing elements within the same group.
V. Factors Affecting Acidity: Resonance Effects
If the conjugate base of an acid is stabilized by resonance, the acid will be stronger. The classic illustration is the comparison between carboxylic acids (pKa approximately 5) and alcohols (pKa approximately 16). When a carboxylic acid loses a proton, the resulting carboxylate anion distributes its negative charge equally over two oxygen atoms through two equivalent resonance structures. An alkoxide ion, by contrast, has no resonance stabilization and bears its full negative charge on a single oxygen atom.
A similar argument explains why phenol (pKa approximately 10) is more acidic than cyclohexanol (pKa approximately 16). The phenoxide anion delocalizes its negative charge into the aromatic ring through four resonance structures, while the cyclohexanol anion cannot achieve any such delocalization. The same principle applies to amides versus amines: the N-H bond of an amide (pKa approximately 25) is more acidic than that of an amine (pKa approximately 38) because the amidate anion is resonance-stabilized through interaction with the adjacent carbonyl. In general, the more resonance structures that stabilize a conjugate base, the stronger the corresponding acid.
<image>Panel A: Resonance structures of the acetate anion (CH3COO-) showing the negative charge delocalized equally over both oxygen atoms, compared to ethoxide (CH3CH2O-) with the charge localized on one oxygen. A dotted-line hybrid structure of acetate is also shown. Panel B: Resonance structures of the phenoxide anion showing delocalization of the negative charge from oxygen into the ortho and para positions of the aromatic ring, with all four contributing structures drawn and curved arrows showing electron movement.</image>
VI. Factors Affecting Acidity: Inductive Effects
The inductive effect refers to the pull or push of electron density through sigma bonds by electronegative or electropositive substituents. Electron-withdrawing groups increase acidity because they stabilize the conjugate base by drawing electron density away from the negative charge. For example, chloroacetic acid (pKa 2.87) is substantially more acidic than acetic acid (pKa 4.76), and trichloroacetic acid (pKa 0.64) is even stronger due to the cumulative effect of three chlorine atoms. Other electron-withdrawing groups include nitro groups, cyano groups, and carbonyl groups.
Inductive effects diminish with distance from the acidic site. This is demonstrated by comparing 2-chlorobutanoic acid (pKa 2.86), 3-chlorobutanoic acid (pKa 4.05), and 4-chlorobutanoic acid (pKa 4.52), which show a clear decrease in the chlorine's effect as the number of intervening bonds increases. Electron-donating groups, such as alkyl substituents, have the opposite effect: they destabilize the conjugate base and reduce acidity. When both inductive and resonance effects are present in the same molecule, resonance effects are generally the stronger influence.
VII. Factors Affecting Acidity: Hybridization Effects
The hybridization of a carbon atom has a profound effect on the acidity of its C-H bonds. Terminal alkynes with sp-hybridized carbon have the most acidic C-H bonds (pKa approximately 25), followed by sp2 C-H bonds in vinyl and aryl compounds (pKa approximately 44), and finally sp3 C-H bonds in alkanes (pKa approximately 50), which are the least acidic.
The explanation lies in the s character of the orbital holding the electron pair of the conjugate base. An sp orbital has 50% s character, meaning electrons in this orbital are held closer to the carbon nucleus and are therefore more stabilized. An sp2 orbital has 33% s character, and an sp3 orbital has only 25%. Greater s character effectively increases the electronegativity of the carbon, stabilizing the resulting anion and making the parent acid stronger. This has practical importance: terminal alkynes can be quantitatively deprotonated by bases like sodium amide (NaNH2), whose conjugate acid (NH3, pKa approximately 38) is far weaker than the alkyne.
VIII. Predicting Acid-Base Reactions in Organic Chemistry
To predict whether an acid-base reaction will proceed, one compares the pKa of the acid on each side of the equation. Equilibrium favors the side with the weaker acid (higher pKa). A practical rule follows: a base can deprotonate an acid only if the pKa of that acid is lower than the pKa of the base's conjugate acid.
Several common bases in organic chemistry illustrate this principle. Sodium hydroxide (NaOH) has water as its conjugate acid (pKa 15.7), so it can deprotonate acids with pKa below about 15. Sodium hydride (NaH), with conjugate acid H2 (pKa 35), and sodium amide (NaNH2), with conjugate acid NH3 (pKa 38), are much stronger bases capable of deprotonating a wider range of substrates. Butyllithium (n-BuLi), with conjugate acid butane (pKa approximately 50), is among the strongest bases available. LDA (lithium diisopropylamide), with conjugate acid diisopropylamine (pKa approximately 36), is another widely used strong base.
Consider whether NaOH can deprotonate ethanol. The pKa of ethanol is 16, while the pKa of water is 15.7. Because these values are nearly identical, the equilibrium slightly favors the reactants, and NaOH does not fully deprotonate ethanol in practice.
<image>A reaction coordinate diagram comparing two acid-base reactions. Reaction 1: acetic acid (pKa 4.76) + NaOH showing favorable equilibrium (Keq >> 1) with the product side lower in energy. Reaction 2: ethanol (pKa 16) + NaOH showing an unfavorable equilibrium (Keq < 1) with the reactant side lower in energy. Energy levels are labeled with the pKa values of the acids on each side. Arrows indicate which direction equilibrium favors in each case.</image>
IX. Basicity in Organic Chemistry
Basicity is simply the reverse perspective of acidity. A strong base has a weak conjugate acid (high pKa), while a weak base has a strong conjugate acid (low pKa). Any factor that stabilizes a negative charge makes a species less basic and its conjugate acid more acidic. Common organic bases can be ranked by strength: weak bases include carboxylate anions, water, and alcohols; moderate bases include amines and alkoxides; and strong bases include hydride (H-), amide (NH2-), and carbanions (R-).
It is important to distinguish basicity from nucleophilicity. Basicity is a thermodynamic property that measures affinity for a proton, while nucleophilicity is a kinetic property that measures affinity for an electrophilic carbon. Although the two often correlate, they are not identical, and this distinction becomes crucial in the study of substitution reactions in later lectures.
