Chapter 1: Structure and Bonding (McMurry Organic Chemistry, 9th Edition)

Practice flashcards covering core concepts from Chapter 1: Structure and Bonding, including atomic structure, orbitals, electron configurations, bonding theories, hybridization, and basic MO theory.

What is organic chemistry as described in the notes?

The study of carbon compounds; organisms are made of organic chemicals such as proteins, DNA, foods, and medicines.

Who demonstrated that an inorganic salt could be converted into an organic substance (urea) in 1828?

Friedrich Wöhler.

What key 1816 discovery by Chevreul is mentioned in origins of organic chemistry?

Soap can be separated into several organic compounds called fatty acids.

What group and how many valence electrons does carbon have, and what does that enable it to do?

Carbon is a group 4A element with four valence electrons, enabling it to form four covalent bonds.

What is the approximate diameter of an atom?

About 2 × 10^-10 meters (approximately 200 picometers).

What do atomic number Z and mass number A represent?

Z is the number of protons; A is protons plus neutrons.

What are isotopes?

Atoms of the same element with the same atomic number (Z) but different mass numbers (A).

What is an orbital in atomic structure?

A region described by a wave function (ψ) where an electron is likely to be found; described by ψ and ψ^2.

Which orbitals are most common in organic and biological chemistry?

s and p orbitals.

Describe the shapes of s and p orbitals.

s orbitals are spherical; p orbitals are dumbbell-shaped with lobes and a node.

How are orbitals organized in shells, and how many electrons can each orbital hold?

Orbitals are grouped into shells around the nucleus; each orbital can hold two electrons; shells have increasing size and energy.

What are the three p orbitals, and how are they designated?

px, py, and pz, mutually perpendicular.

What is Hund’s rule?

Electrons occupy degenerate orbitals singly with parallel spins before pairing.

What is the Aufbau principle?

Electrons fill the lowest-energy orbitals first, in the order 1s < 2s < 2p < 3s < 3p < 4s < 3d.

What is the ground-state electron configuration of sulfur (S)?

1s^2 2s^2 2p^6 3s^2 3p^4.

What does the Pauli exclusion principle state about electrons in an orbital?

No more than two electrons per orbital, and they must have opposite spins.

What is the bond energy released when two hydrogen atoms form H2, and what are the energy units?

436 kJ/mol; convert using 1 kJ = 0.2390 kcal (1 kcal = 4.184 kJ).

What is a valence bond, in simple terms?

A covalent bond formed by the overlap of singly occupied atomic orbitals.

What is a sigma (σ) bond?

A bond formed by head-on overlap with cylindrical symmetry.

What is the relationship between bond length and bond strength?

There is an optimal bond length for maximum stability; too close leads to repulsion, too far weakens bonding.

Describe sp3 hybridization and name a representative molecule.

Mixing of one s and three p orbitals to form four equivalent sp3 orbitals; exemplified by methane, CH4.

What are the C–H bond length and bond energy in methane (CH4)?

C–H bond length ≈ 109 pm; bond strength ≈ 439 kJ/mol.

What is the structural takeaway for ethane (C2H6) in terms of bonds and angles?

Two sp3 carbons connected by a C–C σ bond; six C–H bonds; C–H ~421 kJ/mol; C–C ≈154 pm; bond angles ≈109° (tetrahedral).

What is sp2 hybridization and how does it relate to ethylene (C2H4)?

sp2 = one s and two p combine to form three in-plane sp2 orbitals; one unhybridized p orbital remains for π bonding, giving planar geometry with ~120° angles and a C=C double bond.

Describe the bonding in ethylene (C2H4).

σ bond from sp2–sp2 overlap and a π bond from sideways overlap of unhybridized p orbitals, forming a C=C double bond.

How does formaldehyde illustrate sp2 hybridization of carbon?

Carbon is sp2-hybridized, forming three sp2 orbitals for σ bonds and one unhybridized p orbital for the π bond with oxygen.

What is acetylene’s (C2H2) hybridization and geometry?

Each carbon is sp-hybridized; linear geometry (180°); sp–sp σ bond between carbons and two π bonds from unhybridized p orbitals.

Differentiate sigma (σ) and pi (π) bonds.

σ bonds result from end-on overlap with cylindrical symmetry; π bonds arise from sideways overlap of parallel p orbitals.

How does sulfur in dimethyl sulfide (CH3)2S relate to hybridization?

S is described as approximately sp3-hybridized with significant deviation from ideal tetrahedral angles due to lone pairs.

What is the role of lone pairs in nonbonding electrons?

Lone pairs are valence electrons not used in bonding; e.g., nitrogen in NH3 has one lone pair.

What is the hybridization and geometry of oxygen in methanol (CH3OH)?

Oxygen is sp3-hybridized; C–O–H bond angle ≈ 108.5°; two lone pairs on O occupy two sp3 orbitals.

What is the typical O–P–O bond angle range in methyl phosphate and its implication for phosphorus hybridization?

Approximately 110° to 112°; implies sp3-like hybridization around phosphorus.

What does molecular orbital (MO) theory describe?

Covalent bonds as a result of the combination of atomic orbitals to form molecular orbitals; bonding MOs are lower in energy, antibonding MOs are higher.

In MO theory, which molecular orbital is occupied for H2?

The bonding σ MO is occupied; the antibonding MO remains unoccupied.

What are condensed and skeletal structures, and how do they differ?

Condensed structures show bonds implicitly; skeletal structures hide carbon atoms at line junctions, with hydrogens on carbon often omitted; non-C, H atoms are shown.

What are Kekulé structures and how do they relate to skeletal structures?

Kekulé (line-bond) structures explicitly show bonds between atoms; skeletal structures summarize the carbon framework with bonds not always drawn.

Give examples of compounds used to illustrate Kekulé vs skeletal representations (as per the notes).

Isoprene, methylcyclohexane, and phenol are shown to compare Kekulé and skeletal drawings.

What rule helps determine how many hydrogens are bonded to each carbon in organic molecules?

Carbons typically form four bonds total (tetravalence); hydrogens fill to complete an octet as needed, following typical valence patterns.

What is the general takeaway about carbon’s bonding modes in organic chemistry from the summary?

Carbon forms single bonds using sp3 hybrids (tetrahedral), double bonds using sp2 hybrids (planar), and triple bonds using sp hybrids (linear).