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CHEM 2323: Chapter 2 - Acids, Bases, and Functional Groups

Chapter 2: Acids and Bases; Functional Groups

  • Examples of complex acids discussed include citric acid and ascorbic acid, highlighting the presence of multiple hydroxyl (OH) and carboxyl (COOH) groups.

Bond Dipole Moment

  • Bond polarities exist on a spectrum from nonpolar covalent, through polar covalent, to totally ionic.

  • Dependence: Bond polarity is influenced by the amount of charge separation and the distance between the charges.

  • Measurement: Bond dipole moments (\mu) are measured in debyes (D).

Bond Dipole Moments (Debye) for Some Common Covalent Bonds

Bond

Dipole Moment, (\mu)

Bond

Dipole Moment, (\mu)

C-N

0.22 D

H-C

0.3 D

C-O

0.86 D

H-N

1.31 D

C-F

1.51 D

H-O

1.53 D

C-Cl

1.56 D

C=O

2.4 D

C-Br

1.48 D

C\equivN

3.6 D

C-I

1.29 D

Molecular Dipole Moment

  • The molecular dipole moment is the vector sum of individual bond dipole moments.

  • Dependence: It depends on both bond polarity and molecular geometry (bond angles).

  • Contribution of Lone Pairs: Lone pairs of electrons significantly contribute to the overall molecular dipole moment.

Effects of Lone Pairs on Dipole Moments (Examples)

  • Ammonia (NH_3): (\mu = 1.5D)

  • Water (H_2O): (\mu = 1.9D)

  • Acetone (CH3COCH3): (\mu = 2.9D)

  • Acetonitrile (CH_3CN): (\mu = 3.9D)

  • Self-Exercise: Draw the bond and molecular dipoles for water (H2O) and carbon dioxide (O=C=O). (Note: CO2 has polar bonds but a nonpolar molecule due to symmetrical geometry, resulting in a zero net molecular dipole moment.)

Intermolecular Forces

  • The strength of attractions between molecules dictates crucial physical properties such as melting point (m.p.), boiling point (b.p.), and solubility of compounds.

  • Classification of Attractive Forces:

    • Dipole–dipole forces: Occur between polar molecules and are measured in debyes.

    • London dispersion forces: Weak, temporary attractive forces, present in all molecules but dominant in nonpolar ones.

    • Hydrogen bonding: A strong type of dipole-dipole interaction specific to molecules containing —OH or —NH groups.

Dipole-Dipole Interactions - Attraction

  • These are common attractions between polar molecules.

  • Symbolized by (\delta+) and (\delta-) to indicate partial positive and negative charges, respectively.

  • Example: In chloromethane (CH_3Cl), the carbon-chlorine bond is polar, with chlorine being (\delta-) and carbon (\delta+). Molecules align such that the (\delta+) region of one molecule attracts the (\delta-) region of another.

  • Self-Exercise: Which of the two isomers, cis-1,2-dichloroethene or trans-1,2-dichloroethene, has a higher boiling point? (Hint: Consider the molecular dipole moments and resulting dipole-dipole interactions.)

London Dispersion Forces

  • One of the Van der Waals forces, these are the main attractive forces in nonpolar molecules.

  • Mechanism: A temporary, instantaneous attractive dipole–dipole interaction that arises for a fraction of a second due to momentary distortions in electron distribution.

  • Proportionality: This attractive force is roughly proportional to the molecular surface area. Larger surface area allows for more points of contact and temporary dipole formation.

  • Polarizability: Larger atoms possess more loosely held electrons and are, therefore, more polarizable, leading to stronger London dispersion forces.

Effect of Branching on Boiling Point

  • n-Pentane (linear isomer): Has the greatest surface area, leading to stronger London dispersion forces and thus the highest boiling point.

  • Increased Branching: As the amount of chain branching increases, the molecule's shape becomes more spherical, and its surface area decreases.

  • Neopentane (most highly branched isomer): Has the smallest surface area, resulting in the weakest London dispersion forces and the lowest boiling point among its isomers.

Hydrogen Bonding

  • Hydrogen bonding is a particularly strong dipole–dipole attraction.

  • Requirements: Organic molecules must possess N—H or O—H groups to be able to form a hydrogen bond. The hydrogen atom involved is bonded to a highly electronegative atom (N or O).

  • Mechanism: The hydrogen from one molecule, being (\delta+), is strongly attracted to a lone pair of electrons on the oxygen or nitrogen of an adjacent molecule.

  • Strength Difference: The O—H bond is more polar than the N—H bond, making alcohols capable of stronger hydrogen bonding compared to amines.

How Intermolecular Forces Affect Boiling Points

  • Hydrogen Bonding and Boiling Point: The presence of hydrogen bonding significantly increases the boiling point of a molecule.

  • Alcohols vs. Amines: Due to the higher polarity of the O—H bond compared to the N—H bond, alcohols exhibit stronger hydrogen bonding than amines and consequently have higher boiling points.

  • Self-Exercise: Rank compounds A, B, C, D, E by boiling point (1 having the highest, 5 the lowest). For more problems, visit: https://chemistry.utdallas.edu/ochemrank/

Polarity Effects on Solubility

  • General Rule: "Like dissolves like."

    • Polar solutes: Dissolve readily in polar solvents.

    • Nonpolar solutes: Dissolve readily in nonpolar solvents.

  • Principle: Molecules with similar types and strengths of intermolecular forces will mix freely with each other.

Polar Solute in Polar Solvent

  • When a polar solute is placed in a polar solvent, the attractive forces between solute and solvent molecules can overcome the solute-solute interactions and solvent-solvent interactions.

  • Energy and Entropy: This process often involves the release of energy (hydration) and an increase in entropy, favoring dissolution.

Polar Solute in Nonpolar Solvent

  • In this scenario, the nonpolar solvent's weak intermolecular attractions are insufficient to break apart the strong intermolecular interactions within the polar solute.

  • Outcome: The polar solid will generally not dissolve in the nonpolar solvent, as the energy cost for breaking the polar solute's interactions is too high.

Nonpolar Solute in Nonpolar Solvent

  • The weak intermolecular attractions characteristic of a nonpolar substance can be readily overcome by interaction with a nonpolar solvent, which also has weak attractions.

  • Outcome: The nonpolar substance dissolves, as the energy change is favorable or negligible, and entropy often increases.

Nonpolar Solute with Polar Solvent

  • If a nonpolar molecule were to attempt to dissolve in water (a polar solvent), it would disrupt the hydrogen bonds present between water molecules.

  • Outcome: This disruption is energetically unfavorable, meaning nonpolar substances generally do not dissolve in water.

Hydrophobic and Hydrophilic

  • Hydrophobic ("water-hating"): Nonpolar substances that cannot readily dissolve in water are termed hydrophobic. Example: Motor oil and water do not mix.

  • Hydrophilic ("water-loving"): Polar substances that dissolve easily in water are termed hydrophilic.

Arrhenius Acids and Bases

  • Arrhenius Acids: Substances that dissociate in water (H2O) to yield hydronium ions (H3O^+).

    • Strength: Stronger acids dissociate to a greater extent than weaker acids.

  • Arrhenius Bases: Substances that dissociate in water to yield hydroxide ions (OH^-).

    • Strength: Stronger bases (e.g., NaOH) dissociate more completely than weaker bases (e.g., Mg(OH)_2).

Brønsted-Lowry Acids and Bases

  • Brønsted–Lowry Acids: Defined as any chemical species that can donate a proton (H^+).

  • Brønsted–Lowry Bases: Defined as any chemical species that can accept a proton (H^+).

Lewis Acid–Base Theory

  • Lewis Acids: Any species that can accept a pair of electrons. They typically possess an empty orbital and are often referred to as electrophiles (electron-loving).

  • Lewis Bases: Any species that can donate a lone pair of electrons. They are often referred to as nucleophiles (nucleus-loving) due to their attraction to positive centers.

    • Diagram: An empty orbital can accept electrons (Lewis Acid); a lone pair can donate electrons (Lewis Base).

Conjugate Acids and Bases

  • Conjugate Acid: When a Brønsted-Lowry base accepts a proton, it forms its conjugate acid, which is now capable of donating that proton back.

  • Conjugate Base: When a Brønsted-Lowry acid donates its proton, it forms its conjugate base, which is now capable of accepting that proton back.

Acid Strength

  • The strength of a Brønsted–Lowry acid is quantified by the extent of its ionization in water.

  • Acid-Dissociation Constant (Ka): The equilibrium constant for the ionization of an acid, Ka, indicates the relative strength of the acid. A stronger acid has a larger K_a value.

  • pKa Value: Because Ka values often span a wide numerical range, they are commonly expressed in logarithmic form as ( pKa = -log(Ka)). A weaker acid has a larger pK_a value (i.e., less negative or more positive).

Base Strength

  • The strength of a base is measured using the equilibrium constant (K_b) of its hydrolysis reaction (reaction with water).

  • Base-Dissociation Constant (Kb): A stronger base has a larger Kb value.

  • pKb Value: Similar to acids, base strength is often expressed as (pKb = -log(Kb)). A weaker base has a larger pKb value.

Relative Strength of Some Common Acids and Their Conjugate Bases

Acid

Conjugate Base

(K_a)

(pK_a)

HCl

Cl^-

1 \times 10^7

-7

H_3O^+

H_2O

55.6

-1.7

HF

F^-

6.8 \times 10^{-4}

3.17

Formic acid

Formate ion

1.7 \times 10^{-4}

3.76

Acetic acid

Acetate ion

1.8 \times 10^{-5}

4.74

Hydrocyanic acid

Cyanide ion

6.0 \times 10^{-10}

9.22

NH_4^+

NH_3

5.8 \times 10^{-10}

9.24

Water

Hydroxide ion

1.8 \times 10^{-16}

15.7

Ethyl alcohol

Ethoxide ion

1.3 \times 10^{-16}

15.9

Ammonia

Amide ion

10^{-36}

36

Methane

Methyl anion

<10^{-50}

50

  • Trend: As acid strength increases (larger Ka, smaller pKa), the strength of its conjugate base decreases. Conversely, a weaker acid has a stronger conjugate base. This relationship is always inverse.

Equilibrium Positions of Acid–Base Reactions

  1. Favorability: Acid–base equilibrium always favors the formation of the weaker acid and the weaker base.

  2. pKa/pKb Values: The weaker acid will have the larger pKa value. The weaker base will have the larger pKb value.

  3. Side of Equation: The weaker acid and the weaker base will always be on the same side of the chemical equation (either both as reactants or both as products).

  4. Prediction: The favored side of the equilibrium can be predicted by comparing the strengths (or pK_a values) of the two acids involved or the two bases involved in the reaction.

Solvent Effects on Acidity and Basicity

  • Water as Amphoteric Solvent: Water (H_2O) is an amphoteric substance, meaning it can act as both an acid and a base.

    • Its conjugate acid is the hydronium ion (H3O^+), with a (pKa = -1.7).

    • Its conjugate base is the hydroxide ion (OH^-), with a (pK_b = -1.7).

  • Leveling Effect: Any acid stronger than H3O^+ (i.e., with a pKa more negative than -1.7) will be completely deprotonated by water, effectively becoming H3O^+. Similarly, any base stronger than OH^- (i.e., with a pKb more negative than -1.7) will be completely protonated by water, effectively becoming OH^-. This means water