Organic Chemistry Structure, Bonding, and Acid-Base Chemistry Syllabus

Structure, Bonding, and Acid-Base Chemistry syllabus — full chapter and topic list

Expand any chapter to see its topics and sub-topics. This is the whole examinable outline for Structure, Bonding, and Acid-Base Chemistry in Organic Chemistry, not a summary of it.

1. 1

   Atomic Structure and Chemical Bonding

   5 topics

   • Atomic Orbitals and Electron Configuration
   • Ionic and Covalent Bonds
   • Lewis Structures
   • Hybridization
   • Resonance and Delocalization
2. 2

   Molecular Representations and Geometry

   4 topics

   • Condensed and Skeletal Structures
   • VSEPR Theory and Molecular Shape
   • Bond Angles and Bond Lengths
   • Degrees of Unsaturation
3. 3

   Functional Groups and Nomenclature

   4 topics

   • Identifying Functional Groups
   • IUPAC Naming Rules
   • Naming Common Functional Classes
   • Constitutional Isomers
4. 4

   Acids and Bases in Organic Chemistry

   5 topics

   • Bronsted-Lowry Acids and Bases
   • pKa and Acid Strength
   • Lewis Acids and Bases
   • Nucleophiles and Electrophiles
   • Curved Arrow Notation
5. 5

   Intermolecular Forces and Physical Properties

   3 topics

   • London Dispersion Forces
   • Dipole-Dipole and Hydrogen Bonding
   • Boiling Point, Melting Point, and Solubility

Structure, Bonding, and Acid-Base Chemistry flashcards for Organic Chemistry

21 of 50 cards from the Structure, Bonding, and Acid-Base Chemistry deck — real questions with worked answers.

1. What is an atomic orbital, and how many electrons can a single orbital hold?

   An atomic orbital is a region of space around a nucleus where there is a high probability of finding an electron (a solution to the Schrödinger equation). Each orbital holds a maximum of 2 electrons with opposite spins (Pauli exclusion principle).

2. State the electron configuration of a ground-state carbon atom and give its number of valence electrons.

   Carbon ($Z=6$): $1s^{2}\,2s^{2}\,2p^{2}$. It has 4 valence electrons (in the $n=2$ shell).

3. State Hund's rule and the Aufbau principle for filling atomic orbitals.

   Aufbau principle: electrons fill the lowest-energy orbitals first. Hund's rule: within a set of degenerate orbitals, electrons occupy separate orbitals singly with parallel spins before pairing up.

4. How do the shapes of $s$ and $p$ atomic orbitals differ?

   An $s$ orbital is spherical and centered on the nucleus. A $p$ orbital is dumbbell-shaped (two lobes) with a node at the nucleus; the three $p$ orbitals ($p_x$, $p_y$, $p_z$) are oriented along perpendicular axes.

5. What distinguishes an ionic bond from a covalent bond?

   An ionic bond forms by transfer of electrons, producing oppositely charged ions held by electrostatic attraction (large electronegativity difference, typically $\Delta EN > 1.7$). A covalent bond forms by sharing of electron pairs between atoms (smaller $\Delta EN$).

6. What is the difference between a polar covalent and a nonpolar covalent bond?

   A nonpolar covalent bond shares electrons equally ($\Delta EN \approx 0$, e.g. $\ce{C-C}$). A polar covalent bond shares electrons unequally, creating partial charges ($\delta+$ and $\delta-$), when $0 < \Delta EN < 1.7$ (e.g. $\ce{C-O}$).

7. State the octet rule and name two common exceptions among second-row and beyond elements.

   The octet rule states that atoms tend to gain, lose, or share electrons to achieve 8 valence electrons (a full $s$ and $p$ shell). Exceptions: hydrogen (wants 2, a duet), boron/beryllium (often electron-deficient, 6 or fewer), and elements in period 3+ like phosphorus and sulfur that can expand their octet.

8. Give the step-by-step procedure for drawing a Lewis structure.

   1) Count total valence electrons (add for negative charge, subtract for positive). 2) Place the least electronegative atom in the center (never H). 3) Connect atoms with single bonds. 4) Distribute remaining electrons as lone pairs to complete octets on outer atoms first. 5) Form multiple bonds if the central atom lacks an octet. 6) Assign formal charges to check the best structure.

9. What is the formula for formal charge, and use it for the oxygen in the hydronium ion $\ce{H3O+}$.

   $$\text{Formal charge} = (\text{valence } e^-) - (\text{nonbonding } e^-) - \tfrac{1}{2}(\text{bonding } e^-)$$ For O in $\ce{H3O+}$: $6 - 2 - \tfrac{1}{2}(6) = 6 - 2 - 3 = +1$.

10. What are the hybridization states $sp$, $sp^{2}$, and $sp^{3}$, and how many hybrid orbitals does each produce?

    $sp$: one $s$ + one $p$ → 2 hybrid orbitals (linear). $sp^{2}$: one $s$ + two $p$ → 3 hybrid orbitals (trigonal planar). $sp^{3}$: one $s$ + three $p$ → 4 hybrid orbitals (tetrahedral).

11. How can you determine the hybridization of an atom by counting groups?

    Count the number of sigma bonds plus lone pairs (the steric number) on the atom. Steric number 2 → $sp$; 3 → $sp^{2}$; 4 → $sp^{3}$.

12. What is the hybridization and geometry of each carbon in ethene ($\ce{C2H4}$) and ethyne ($\ce{C2H2}$)?

    Ethene: each carbon is $sp^{2}$, trigonal planar (double bond). Ethyne: each carbon is $sp$, linear (triple bond).

13. Describe the difference between a sigma ($\sigma$) bond and a pi ($\pi$) bond.

    A $\sigma$ bond forms by head-on (end-to-end) overlap of orbitals along the internuclear axis and allows free rotation. A $\pi$ bond forms by side-by-side overlap of $p$ orbitals above and below the axis, is weaker, and prevents rotation. A double bond = 1 $\sigma$ + 1 $\pi$; a triple bond = 1 $\sigma$ + 2 $\pi$.

14. What is resonance, and how do resonance structures relate to the real molecule?

    Resonance is the delocalization of electrons represented by two or more valid Lewis structures (resonance contributors) that differ only in electron placement. The real molecule is a resonance hybrid — a single weighted average of all contributors, not any one structure alternating.

15. List three rules for identifying the major (most stable) resonance contributor.

    1) Maximize the number of full octets (especially on electronegative atoms). 2) Minimize formal charges and separation of opposite charges. 3) Place negative charge on the more electronegative atom and positive charge on the less electronegative atom.

16. Which electrons can participate in resonance delocalization?

    Only $\pi$ electrons (in double/triple bonds), lone pairs adjacent to a $\pi$ system or empty orbital, and electrons in $p$ orbitals. Sigma bonds and lone pairs isolated from $\pi$ systems do not delocalize by resonance.

17. What is the difference between a condensed structural formula and a skeletal (line-angle) structure?

    A condensed formula writes atoms in sequence and omits most bond lines, grouping H's with their carbon (e.g. $\ce{CH3CH2OH}$). A skeletal structure shows carbon skeletons as lines where each vertex/endpoint is a carbon; carbons and their attached hydrogens are implied, and heteroatoms are drawn explicitly.

18. In a skeletal structure, how do you determine the number of hydrogens on a given carbon?

    Each neutral carbon forms 4 bonds. Count the bonds shown to that carbon (including multiple bonds and bonds to heteroatoms); the number of implied hydrogens = $4 - (\text{number of bonds shown})$.

19. State the main postulate of VSEPR theory.

    Valence Shell Electron Pair Repulsion theory states that electron groups (bonds and lone pairs) around a central atom arrange themselves as far apart as possible to minimize electron-electron repulsion, which determines molecular geometry.

20. Give the electron-group geometry and ideal bond angle for steric numbers 2, 3, and 4.

    Steric number 2: linear, $180^{\circ}$. Steric number 3: trigonal planar, $120^{\circ}$. Steric number 4: tetrahedral, $109.5^{\circ}$.

21. How does a lone pair affect the molecular shape and bond angles compared with the electron-group geometry? Illustrate with water.

    Lone pairs repel more strongly than bonding pairs, so they compress bond angles and the molecular shape (which counts only atoms) differs from electron geometry. Water ($\ce{H2O}$) has 4 electron groups (tetrahedral electron geometry) but a bent molecular shape with an angle of about $104.5^{\circ}$ (less than $109.5^{\circ}$).

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