CHM 221 Organic Chemistry I Chapter 1 Review

Flashcards covering key concepts from Chapter 1 review of Organic Chemistry I, including structure, bonding, orbitals, and intermolecular forces.

What is Organic Chemistry?

Chemistry of carbon-containing compounds; organic compounds are built around carbon–carbon bonds.

What did the Vital Force theory claim and who disproved it?

The theory claimed organic compounds could only arise from living organisms; Friedrich Wöhler disproved it in 1828 by synthesizing urea from inorganic reagents.

Why is the molecular formula C2H6O not sufficient to define a molecule?

Because substances are defined by structure; compounds with the same formula but different structures are structural (constitutional) isomers.

What is the HONC 1234 rule?

Hydrogen forms 1 bond; Oxygen forms 2 bonds; Nitrogen forms 3 bonds; Carbon forms 4 bonds; Halogens typically form 1 bond.

What is the nucleus in an atom?

The central region containing protons and neutrons.

What are electrons in atomic structure?

Negatively charged particles that orbit the nucleus in regions called orbitals.

What does the atomic number represent?

The number of protons in the nucleus.

What does the mass number represent?

The total number of protons and neutrons in the nucleus.

What are core electrons?

Electrons close to the nucleus that are not generally involved in bonding.

What are valence electrons?

Electrons in the outermost shell; counted by group on the periodic table for main-group elements.

What are simple Lewis structures and how are they built?

Draw atoms with valence electrons as dots, connect atoms with bonds, ensure each atom has an octet, and pair unpaired electrons as needed (H and halogens often form 1 bond and are placed last).

What is formal charge?

Deviation from the expected valence count; FC can be calculated as valence electrons minus half of bonding electrons minus nonbonding electrons.

What are partial charges and electronegativity?

Atoms in molecules can have partial charges due to uneven electron distribution; electronegativity is the tendency of an atom to attract electrons in a bond.

What are the three types of bonds and their key characteristics?

1) Covalent: electrons shared between two atoms; 2) Polar covalent: electrons shared but unequally due to differing electronegativities; 3) Ionic: electrons transferred to achieve full valence shells.

What is the Aufbau principle?

Electrons fill the lowest available energy orbitals first.

What is the Pauli exclusion principle?

No two electrons in an atom can have the same set of quantum numbers; each orbital holds at most two electrons with opposite spins.

What is Hund’s rule?

For degenerate orbitals, electrons occupy separate orbitals with parallel spins before pairing up.

What are the shapes of s and p atomic orbitals?

s orbitals are spherical; p orbitals are dumbbell-shaped.

What is the significance of the phases of atomic orbitals?

The sign (phase) of orbitals affects constructive vs. destructive overlap when forming bonds.

What is a sigma (σ) bond and how is it formed?

A sigma bond forms from head-to-head overlap of orbitals (e.g., s-s or sp3-sp3) along the bond axis.

What is a pi (π) bond and how is it formed?

A pi bond forms from side-to-side overlap of unhybridized p orbitals above and below the bond axis.

What is VSEPR theory and what does steric number determine?

VSEPR determines molecular geometry; steric number is the number of groups (atoms or lone pairs) around a central atom.

What are the typical hybridizations and bond angles for 2, 3, and 4 groups around an atom?

2 groups: sp, 180°; 3 groups: sp2, 120°; 4 groups: sp3, 109.5°.

What is the sp3 molecular geometry and what is methane’s geometry?

sp3 corresponds to a tetrahedral geometry with bond angles ~109.5°; CH4 has four equal C–H bonds in a tetrahedral arrangement.

Why is carbon in methane described with sp3 hybridization?

To allow four equivalent bonds with ~109.5° angles; HOwever, hybridization explains equal bond angles not by excitation alone, but by mixing orbitals.

What is the shape and character of sp3 hybrids?

The sp3 hybrid orbital has 25% s-character and 75% p-character, forming four equivalent sp3 orbitals.

What is the hybridization in ethene (C2H4) and what bonds are formed?

Each carbon is sp2-hybridized; one unhybridized p orbital on each carbon forms a π bond, while sp2 orbitals form σ bonds.

What is the hybridization in ethyne (C2H2) and what bonds are formed?

Each carbon is sp-hybridized; unhybridized p orbitals form two π bonds (one on each side) and sp orbitals form σ bonds.

What is Molecular Orbital (MO) Theory in brief?

Atomic orbitals combine to form molecular orbitals; number of MOs equals number of AOs used; bonding MOs are lower in energy, antibonding MOs higher.

What is bond order in MO theory?

Bond order = (number of bonding electrons − number of antibonding electrons) ÷ 2.

What are HOMO and LUMO?

HOMO is the highest occupied molecular orbital; LUMO is the lowest unoccupied molecular orbital; their locations influence reactivity.

How is molecular polarity determined?

By the vector sum of individual bond dipoles; overall dipole moment depends on bond polarity and molecular geometry.

What are the main intermolecular forces?

Dipole-dipole forces, hydrogen bonding, and London dispersion forces.

What is hydrogen bonding and which solvents are protic vs aprotic?

Hydrogen bonding occurs when hydrogen is bonded to N, O, or F; protic solvents can donate H-bonds (e.g., acetic acid); aprotic solvents cannot (e.g., diethyl ether).

What are London dispersion forces and how do they relate to mass and branching?

Weak temporary dipole interactions; heavier molecules have stronger London forces and higher boiling points; more branching reduces surface area, weakening London forces and lowering boiling points.

What is the solubility principle in organic chemistry?

Like dissolves like: polar compounds dissolve in polar solvents; nonpolar compounds dissolve in nonpolar solvents; hydrogen-bonding capability influences mixing.