Organic Chemistry- Exam 1

Ionic bonding

giving up an electron, typically a cation and anion bonded together, however, this will not be common in organic chemistry

Covalent bonds

valence electron sharing

electronegativity definition and trend

• A measure of an atom's attraction for electrons within a bond

• Increases across a period and up a group on the periodic table (Fluorine and Oxygen are the most electronegative

Bond polarity

• Based on electonegativity difference

non polar covalent bond

• electrons are shared relatively equally

• electro negativity difference of less than or equal to 0.4

Polar covalent bond

• Electrons are shared unequally. The more electronegative atom pulls electron density, resulting in a partial negative charge (( \delta^- )), while the other atom acquires a partial positive charge (( \delta^+ )). This is represented by a bond dipole arrow pointing towards the more electronegative atom (e.g., in C-F, Fluorine is ( \delta^- ) and Carbon is ( \delta^+ )).

• Difference in electronegativity of 0.4 to 2

• permanent dipole moment

Ionic bond

An electron is essentially transferred from the less electronegative atom to the more electronegative atom

Electronegativity difference of greater than 2

Formal charge

A charge assigned to an individual atom within a Lewis structure, used to evaluate the best Lewis structure and predict reactivity.

Rule for determining the formal charge of a compound

• The sum of formal charges on all atoms in a molecule must equal the overall charge of the compound.

• Formal charge= # of valence electrons- (# of bonds+ # of lone electrons)

Isomers

• Compounds with the same molecular formula but different properties.

Structural isomers (constitutional)

• Same molecular formula, but different connectivity (arrangement) of atoms.

• Example: C2H6O can be ethanol (CH3CH2OH) or dimethyl ether (CH3OCH3), which have different properties (e.g., hydrogen bonding, boiling points).

Stereoisomers

• Same connectivity of atoms, but different arrangement of atoms in three-dimensional space.

• Examples include optical and geometric isomers (discussed in later chapters).

Skeletal structure rules

• Carbon atoms are implied at vertices and ends of lines.

• Hydrogens attached to carbons are implied to complete octets.

• Heteroatoms (non-carbon/hydrogen) are explicitly drawn.

• Lone pairs on heteroatoms are often omitted but must be remembered.

Resonance structures

•  Structures that have the same placement (connectivity) of atoms but differ in the arrangement of electrons.

  • Distinct from isomers, as atom connectivity remains unchanged.

Resonance Hybrid

• The true structure of a molecule with resonance, which is an average of all contributing resonance structures. No single resonance structure accurately depicts the actual molecule.

  • Example: Nitrite ion (NO2-). The actual N-O bonds are intermediate between single and double bonds, with electrons delocalized over the entire system.

• characteristics of both contributing structures; certain e- pairs are delocalized (spread out) over 2 or more atoms

Major Contributor

The resonance structure that contributes most significantly to the resonance hybrid, often characterized by favorable formal charge distribution.

What does resonance stabilized refer to?

Molecule with 2 or more resonance structures means that is “resonance stabilized”

Add vesper structures (focus is structures 2-4)

Determining if a molecule has a dipole moment?

• A molecule will have a dipole moment if the bonds are polar and they don’t cancel each other out due to geometry molecular shape

Bronstead lowry Acid

• A species that donates a proton (H+).

Bronsted Lowry Base

• A species that accepts a proton (H+).

  • Requires an available electron pair (lone pair or pi bond) to form a new bond with the proton.

  • Examples: Hydroxide (OH-), Ammonia (NH3), carbon-carbon double bonds (pi bonds).

Conjugate Acid

•  Formed when a base accepts a proton.

Conjugate Base

• Formed when an acid donates a proton.

Curved Arrows in an Acid Base reaction

• An arrow starts from the electron pair of the base and points to the proton being accepted from the acid.

• Another arrow starts from the bond between the proton and the rest of the acid, pointing to the atom that retains the electron pair, forming the conjugate base.

Acid Strength

• The tendency of an acid to donate a proton.

• A larger Ka value indicates a stronger acid and smaller pKa

• a stronger acid means that it’s more dissociated

• A stronger acid has a weaker conjugate base and vise versa

pKa values to remember for H2SO4, Carboxylic acids, H2O, Alcohols (-OH), alkanes

-9, 5, 14, 16, 50

Conjugate acid base pairs

• When an acid donates a proton, it forms its conjugate base.

• When a base accepts a proton, it forms its conjugate acid.

Conjugate base stability

Anything that stabilizes the conjugate base makes it a weaker base, which in turn means its corresponding acid is a stronger acid.

Factors that impact acid strength (4)

1. Element to which the proton is attached

2. Inductive Effects

3. Resonance Stabilization

4. Hybridization

Element to which the proton is attached periodic trend

• Down a Group (e.g., HF, HCl, HBr, HI):

  • Acid strength increases down a group (HI > HBr > HCl > HF).

  • Reason: The H-X bond becomes longer and weaker as atomic size increases, making it easier to break. The larger conjugate base (e.g., I-) can better delocalize the negative charge, making it more stable (weaker base). Electronegativity is not the dominant factor here.

• Across a Period (e.g., CH4, NH3, H2O, HF):

  • Acid strength increases across a period (HF > H2O > NH3 > CH4).

  • Reason: Electronegativity increases across a period. A more electronegative atom (e.g., F) is better able to stabilize the negative charge on the conjugate base, making it a weaker base and the acid stronger.

• Summary: Acid strength increases down a group (due to bond strength/size) and across a period (due to electronegativity).

Inductive Effects

• Mechanism: Electronegative atoms (e.g., halogens) located near the acidic proton can pull electron density through sigma bonds.

• Effect: This electron-withdrawing effect helps to stabilize the negative charge on the conjugate base by distributing it, making the conjugate base weaker and the acid stronger.

Resonance Stabilization

• Mechanism: The negative charge on the conjugate base can be delocalized over multiple atoms through resonance (involving pi bonds and lone pairs).

• Effect: Spreading the negative charge over a larger area significantly stabilizes the conjugate base, making it a weaker base and the corresponding acid a stronger acid.

• Example: Carboxylic acids are much stronger than alcohols because the carboxylate conjugate base is resonance-stabilized, while the alkoxide conjugate base is not.

How to determine the stronger base from the skeletal structure

Want a less stable base

1. A negative charge is less stable than an atom with a neutral charge

2. The (-) is less stable when on the smaller atom or the one that is less electronegative

3. (-) is less stable when it’s not shared by other atoms, across multiple resonance structures, or shared by larger atoms

4. (-) is less stable when not surrounded by other electronegative atoms (induction)

5. (-) is less stable when it appears in hybrid orbitals with less S character

• more s character, lower pKa, strong acid+ weak base= stable base ***

Hybridization- it’s effect on acid strength

• Orbital Effect: The type of hybridization of a carbon atom affects the acidity of a C-H bond.

• Acidity Trend: C-H bonds in triple bonds are most acidic, followed by double bonds, and then single bonds (least acidic).

• pKa: Highest pKa for C-H in single bonds, lowest for C-H in triple bonds.

• Mechanism: This effect is explained by the stability of the conjugate base.

  • Hybridization Types:

    • sp3: 4 electron groups (e.g., C in CH4)

    • sp2: 3 electron groups (e.g., C in C=C)

    • sp: 2 electron groups (e.g., C in C≡C)

  • Percent s-character:

    • sp: 50% s-character

    • sp2: 33% s-character

    • sp3: 25% s-character

  • Rule: Increasing percent s-character leads to a more stable conjugate base, which in turn means a stronger corresponding acid.

  • Reason: The s-orbital is closer to the nucleus. Higher s-character means the lone pair in the conjugate base is held closer to the nucleus, increasing electrostatic stabilization of the negative charge.

Equilibrium Direction for predicting proton transfer reactions

• When an acid and a base react, the equilibrium favors the formation of the weaker acid and weaker base.

• net direction of an acid-base reaction depends on relative acid and base strengths- a reaction will proceed to the greater extent in the direction in which a stronger acid and stronger base form a weaker acid and weaker base

Hybridization- simple

• Increasing acidity means that the carbon atom becomes more stable and lowers the hybridization- pKa decreases with this trend

Process for determining proton transfer in acid base reactions

1. Identify the acid and base in the reactants.

2. Draw the products (conjugate acid and conjugate base) using curved arrows to show proton transfer.

3. Compare the strength of the acid (reactant) to its conjugate acid (product), or the base (reactant) to its conjugate base (product).

4. The side with the weaker acid/base pair is favored at equilibrium.

Stronger Acid Definition

• An acid more willing to donate a proton. The reaction proceeds away from the stronger acid. Forms its weaker conjugate base

Weak Acid Definition

•  An acid less willing to give up its proton. Once formed, it is less likely to react in the reverse direction.

Stronger base definition

better able to accept a proton to form its weaker conjugate acid

Br⁻ + CH₃OH Reaction

• Reactants: Br⁻ (base), CH₃OH (acid).

• Most Acidic Proton: The O-H proton in CH₃OH is more acidic than C-H protons due to oxygen's higher electronegativity.

• Proton Transfer: Br⁻ attacks the O-H proton; the O-H bond electrons move to oxygen.

• Products: HBr (conjugate acid), CH₃O⁻ (conjugate base).

• Acid Strength Comparison: HBr is a stronger acid than CH₃OH. This is because bromine is larger than oxygen, allowing the negative charge in Br⁻ (conjugate base of HBr) to be distributed over a larger volume, making it more stable than CH₃O⁻.

• Equilibrium: Since HBr is the stronger acid (on the product side), the equilibrium lies to the left, favoring the reactants (Kc < 1).

Lewis Acids and Bases definition

• Focuses on electron pair donation and acceptance, not necessarily proton transfer.

• Lewis Acid: An electron pair acceptor.

  • Requires an available orbital to accept electrons.

  • Examples:

    • Proton (H⁺): Has an empty 1s orbital.

    • Species with electron-deficient central atoms: BF₃, AlCl₃ (Group 3 elements often have incomplete octets but stable formal charges).

  • Electrophile: Another term for a Lewis acid (electron-seeking, often abbreviated E⁺).

• Lewis Base: An electron pair donor.

  • Examples: NH₃ (has a lone pair), Br⁻ (has lone pairs).

  • Nucleophile: Another term for a Lewis base (nucleus-seeking, often abbreviated Nu⁻).

Alkanes definition

• Contain only carbon-carbon single bonds.

  • General formula: CnH₂n+₂.

  • Examples: Methane (CH₄), Ethane (C₂H₆).

  • Focus of Chapters 2 and 3 (including 3D structure).

Valence Bond Theory and Hybridization of Alkanes

Confused on this…

How to name alkane compounds

• Suffix: Indicates the functional group (e.g., "-ane" for alkanes).

• Parent: Represents the longest continuous carbon chain, determining the root name (e.g., meth-, eth-, prop-, but-, pent)

  • Note: always use the longest chain and it doesn’t have to be a straight line, just continuous)

• Prefix: Describes the identity and location of substituents (side chains or fragments) attached to the parent chain. (always number(s) used to describe the location of the substituents, if attatched to the 2nd and 4th CH3 group, then the prefix would be 2,4-di… the use of di is due to there being 2 substituents)

• Naming of substituents: Based on the number of carbons in the substituent. Uses the parent name root + "-yl" ending (e.g., methyl, ethyl) to denote it as a substituent covalently bonded to the parent chain.

Constitutional Isomers

Molecules sharing the same molecular formula but differing in the connectivity of their atoms.

Butane as an Isomer

• Two constitutional isomers exist.

  • n-Butane: A straight-chain alkane (CH3CH2CH2CH3), often drawn as a zigzag to represent its tetrahedral (sp3 hybridized) carbons.

  • Isobutane: A branched alkane with three carbons in a row and a methyl group attached to the middle carbon.

Pentane as an Isomer

• n-Pentane: A five-carbon straight chain.

• Isopentane: A four-carbon chain with one methyl branch.

• Neopentane: A three-carbon chain with two methyl branches on the central carbon.

Classification of Carbon Atoms

• Carbon atoms are categorized by the number of other carbon atoms they are directly bonded to.

• Primary (1°): Bonded to one other carbon atom (e.g., the carbons in terminal CH3 groups).

• Secondary (2°): Bonded to two other carbon atoms (e.g., carbons in CH2 groups within a chain).

• Tertiary (3°): Bonded to three other carbon atoms (e.g., the central carbon in isobutane).

• Quaternary (4°): Bonded to four other carbon atoms (e.g., the central carbon in neopentane).

Naming Alkyl Substituents

• Methyl (Me): A one-carbon substituent ((-CH_3)).

• Ethyl (Et): A two-carbon substituent ((-CH_2CH_3)).

• Three-Carbon Substituents (from Propane):

  • n-Propyl: Connected via a primary carbon ((-CH_2CH_2CH_3)).

  • Isopropyl: Connected via the secondary carbon ((-CH(CH_3)_2)).

• Four-Carbon Substituents (from n-Butane):

  • n-Butyl: Connected via a primary carbon ((-CH_2CH_2CH_2CH_3)).

  • sec-Butyl: Connected via a secondary carbon ((-CH(CH_3)CH_2CH_3)).

• The naming of substituents depends on the constitutional isomer they originate from and the specific carbon atom through which they attach to the parent chain.