Chapter 01: Structure and Bonding (Wade 9e) - Study Notes

Chapter 01 Outline: Structure and Bonding

Introduction to Chemistry

• Chemistry is the study of the structure and function of molecules and the rules that govern their interactions.

1.1 Origins of Organic Chemistry

• Organic chemistry is defined as the "chemistry of carbon compounds."

  • It includes the study of molecules, physical properties, and the structure of organic compounds.

  • It encompasses both experimental and computational studies.

• The term "organic" historically means "derived from living organisms."

• Organic molecules contain various subunits that control their reactivity. These subunits are classified as functional groups.

1.2 Principles of Atomic Structure

• Atoms are made up of protons, neutrons, and electrons.

  • Protons and neutrons are found in the nucleus.

  • Electrons are in the surrounding cloud.

• Electrons:

  • Are found in shells. The outermost shell is called the valence shell.

  • Valence electrons can be used to predict the behavior of atoms.

  • Electrons have properties of both particles and waves.

  • Atoms gain or lose valence electrons to obtain a more stable configuration, the same as noble gas elements.

  • The octet rule describes how main group elements generally want $8$ valence electrons.

• Electrons are found in orbitals, which are defined as a $3D$ region of space with a $90\%$ probability of where an electron is likely to exist.

  • We can never determine the exact location of an electron (Heisenberg uncertainty principle); however, we can determine the electron density (probability of finding an electron in a certain part of the orbital).

  • A node is a region where there is no probability of finding an electron.

  • s orbital: Spherical shape, with no planar node (region of zero electron density).

  • The energy of orbitals is quantized.

  • p orbital: Consists of $3$ degenerate orbitals.

    • Each orbital has different directionality.

    • A planar node separates the two lobes.

    • Lobes are color-coded to represent peaks and troughs of wave orbital.

• Isotopes are atoms of the same element with a different number of neutrons.

  • Isotopes are often referenced by their mass number ($protons + neutrons$).

  • Example: $<em>{6}^{12}C$ and $</em>{6}^{14}C$.

• Ions are formed when an atom gains or loses an electron or electrons.

  • Example of gain: $\text{Cl(g)} + \text{e}^- \to \text{Cl}^-(\text{g})$

  • Example of loss: $\text{Na(g)} \to \text{Na}^+(\text{g}) + \text{e}^-$

1.3&6 Bond Formation: The Octet Rule & Electronegativity and Bond Polarity

• Electronegativity is the measure of an element's ability to pull electrons through a bond.

  • It can be used to predict the type of bond between two elements.

  • It is used to predict the direction and strength of a bond dipole.

• Ionic Bonding: Electrons are transferred between two atoms.

  • Occurs when there is a large difference ($>2$) in Pauling electronegativity values.

  • Ionic bonds are not as common in organic molecules.

• Covalent Bonding: An electron pair is shared between two atoms.

  • ALL organic molecules contain covalent bonds.

  • Covalent bonds are directional.

  • Nonpolar covalent bonding: Electron pair is shared evenly between the atoms.

    • This typically occurs when there is a small difference ($<0.5$) in Pauling electronegativity values.

  • Polar covalent bonding: Electron pair is not shared evenly between the atoms.

    • This typically occurs when there is a moderate difference ($0.5-2.0$) in Pauling electronegativity values.

    • These bonds have a bond polarity, which can be shown as a vector or delta symbols ($\delta^+, \delta^-$).

1.4,5&7 Lewis Structures, Multiple Bonding & Formal Charges

• G.N. Lewis stated that an atom is most stable if it has a filled outer shell (isoelectronic with a noble gas).

  • The combination of two hydrogen atoms shows this through sharing electrons, giving each hydrogen an electron configuration equal to helium.

  • Covalent bonds are directional.

• Electrostatic potential maps (EPM) are a graphical depiction of electron distribution.

  • Areas of blue indicate electron-poor regions.

  • Areas of red indicate electron-rich regions.

  • Areas of green indicate even electron distribution.

• Differences in electronegativity give rise to bond polarity.

  • The greater the difference in electronegativity, the greater the bond polarity.

• The uneven electron distribution in a compound containing covalent bonds is measured by a quantity called the dipole moment.

  • It is defined to be the amount of charge separation multiplied by the bond length ($\mu = q \times r$).

• Lewis Structures are used to show how valence electrons are arranged in a molecule.

  • Connectivity or constitution (the order the atoms are connected) must be determined to draw a Lewis structure.

  • Bonding electrons (shared pairs) are shown by a solid line.

  • Nonbonding electrons (nonbonding electrons or lone pair electrons) are shown as a pair of dots.

  • The goal is to obtain an octet for all atoms.

    • Exceptions:

      • Hydrogen (H) and Helium (He) obtain $2$ electrons.

      • Third Period elements and above (e.g., S, P, Cl) are capable of having more than $8$ electrons (expanded octets/hypervalency).

• Steps for drawing Lewis Structures:

  1. Calculate the total number of valence electrons.

  2. Draw the molecule with a reasonable skeleton (e.g., the central atom is often the least electronegative or the one from the left side of Periodic Table).

  3. Add electrons to the molecular skeleton to give as many non-hydrogen atoms a full octet of electrons.

     • Odd electrons (unpaired) are unusual and unstable.

     • Elements at the top right of the Periodic Table are more likely to have lone pairs of electrons.

  4. Calculate formal charge.

     • Formal Charge (F.C.) = (valence \ electrons) - (# \ nonbonding \ electrons) - \frac{1}{2}(# \ bonding \ electrons).

     • The sum of the formal charges on individual atoms must equal the net charge of the ion.

     • Formal charges are shown explicitly next to the atom.

• Multiple bonds have more than two electrons shared between atoms to reach octets (e.g., double bonds share $4$ electrons, triple bonds share $6$ electrons).

• Common bonding patterns (useful for quick checking formal charges):

  • Carbon (C):

    • $3$ bonds, $1$ lone pair electrons, Formal Charge: $-1$

    • $4$ bonds, $0$ lone pair electrons, Formal Charge: $0$

    • $3$ bonds, $0$ lone pair electrons, Formal Charge: $+1$

  • Nitrogen (N):

    • $2$ bonds, $2$ lone pair electrons, Formal Charge: $-1$

    • $3$ bonds, $1$ lone pair electron, Formal Charge: $0$

    • $4$ bonds, $0$ lone pair electrons, Formal Charge: $+1$

  • Oxygen (O):

    • $1$ bond, $3$ lone pair electrons, Formal Charge: $-1$

    • $2$ bonds, $2$ lone pair electrons, Formal Charge: $0$

    • $3$ bonds, $1$ lone pair electron, Formal Charge: $+1$

  • Halogen (X):

    • $1$ bond, $3$ lone pair electrons, Formal Charge: $0$

• Lewis Structure Examples:

  • $H_2O$

  • $NH_3$

  • $\text{NO}_2^-$ (Nitrite ion) example breakdown:

    • Calculate valence electrons: $1 \times 5 (N) + 2 \times 6 (O) + 1 (extra \ e^-) = 5 + 12 + 1 = 18 \ e^-$

    • Total valence electrons: $18 + 1 (for \ anion \ charge) = 24 \ e^-$ (Correction from transcript, should be $5 + 2<em>6 + 1 ( ext{charge}) = 18$ total valence electrons, or as shown in original notes $5 + 2</em>6 + 1 = 18$ + extra electron = $24$.) Re-evaluating based on the notes' calculation: N has $5$, O has $6$. For $NO<em>2^-$, N has 5, O has 6. $5 + 2 \times 6 = 17$. Plus $1$ for the charge makes $18$ valence electrons. The transcript example for $NO</em>2^-$ shows $23+1=24e^-$, with N being central. This implies $N \times 5 + 3 \times 6 = 23$, so it suggests an $NO<em>3$ like molecule with a positive charge. The $NO</em>2^-$ on the image is correctly showing $18$ electrons; the $23+1=24$ looks like a typo related to a different example. I will follow the image example. For $NO_2^-$, total valence electrons = N ($5$) + O ($6 \times 2$) + Charge ($1$) = $18$ electrons.

    • Skeleton: O-N-O (N is central, furthest to left as rule of thumb).

    • Add electrons: Start by placing $2$ electrons in each bond ($4$ total). Distribute remaining $14$ electrons as lone pairs to peripheral atoms (O), then central atom (N). One O will need a double bond to satisfy octet.

    • Possible Resonance Structures: A double bond to one oxygen and a single bond to the other, with lone pairs distributed accordingly.

    • Formal Charges (for one resonance structure): For the structure $:O=N-O:^-$ (with $2$ lone pairs on double bonded O, $1$ lone pair on N, $3$ lone pairs on single bonded O).

      • Double bonded O: $6 \text{ (valence)} - 4 \text{ (nonbonding)} - \frac{1}{2}(4 \text{ (bonding)}) = 6 - 4 - 2 = 0$

      • Nitrogen: $5 \text{ (valence)} - 2 \text{ (nonbonding)} - \frac{1}{2}(6 \text{ (bonding)}) = 5 - 2 - 3 = 0$

      • Single bonded O: $6 \text{ (valence)} - 6 \text{ (nonbonding)} - \frac{1}{2}(2 \text{ (bonding)}) = 6 - 6 - 1 = -1$

1.8 Ionic Structures

• Some organic compounds contain ionic bonds, which are NOT shown with a solid line (like covalent bonds).

• Both ions are drawn close to one another.

• Formal charges must be explicitly written on the atoms.

1.9 Resonance

• A single Lewis structure may not be enough to accurately describe a molecule's true structure.

• Nitromethane ($CH<em>3NO</em>2$) is an example where resonance forms are used.

• Spreading electrons and/or charges over two or more atoms stabilizes the ion/molecule.

• General rule: The more resonance structures that can be written, the more stable (lower energy) the molecule is.

• Rules for Drawing Resonance Forms:

  • Only ELECTRONS (pi electrons or lone pairs) are moved.

  • Arrows connecting resonance forms are double-headed ($\leftrightarrow$) to indicate resonance, not equilibrium.

  • Each resonance form must have the same overall charge.

  • The real structure is a resonance hybrid of all valid resonance forms.

• If two resonance structures are not identical, the molecule they represent is a weighted average of the two.

  • One structure may be more important than the other.

• Resonance forms may have different energies.

  • Forms of different energies will contribute to the real structure differently.

  • The major contributor is the more stable resonance form.

  • The minor contributor is the less stable resonance form.

• Criteria for comparing resonance forms:

  1. Has as many octets as possible for each atom (especially C, N, O, F).

  2. Has as many bonds as possible.

  3. Has as little charge separation as possible.

  4. Has the negative charge on the most electronegative atom (e.g., O, N) or a positive charge on the least electronegative atom (e.g., C, often if required).

1.10 Structural Formulas

• There are multiple ways to depict organic molecules:

  • Lewis structures: Show all atoms, bonds, and lone pairs.

  • Condensed Structures: Connectivity is shown without explicitly drawing all bonds.

    • Example: $CH<em>3CH=CHCH</em>2CH<em>2C(CH</em>3)CH_3$

  • Line-angle drawings (also known as skeletal structures or stick figures):

    • Atoms that are not carbon or hydrogen are shown (e.g., N, O, halogens).

    • Multiple bonds (double or triple) are explicitly shown.

    • Bonds are represented by lines.

    • Carbon atoms are implicitly understood to be at the start/end of a line or where two lines meet.

    • Hydrogen atoms bonded to carbon are assumed to be present in the number required to satisfy carbon's valency (assuming a neutral atom unless a formal charge is indicated).

1.11 Molecular Formulas and Empirical Formulas

• Molecular formula gives the exact number of atoms of each element in one molecule of a compound.

• Empirical formula gives the simplest whole-number ratio of atoms in a compound.

1.12 Wave Properties of Electrons in Orbitals

• Viewing an atom as electrons orbiting around the nucleus, while useful, has limitations.

• Quantum mechanics describes an alternate, more detailed view.

• Small particles, like electrons, possess the character of waves.

• A standing wave vibrates in a fixed location.

• A wave function is a mathematical description of the size, shape, and orientation of an orbital.

• The amplitude of a wave function may be positive or negative (this is not related to electrical charge).

• The node is an area of zero amplitude (i.e., zero probability of finding an electron).

1.13 Molecular Orbitals

• Molecular orbitals (MOs) are formed when combined atomic orbitals (AOs) of two DIFFERENT atoms interact during bonding.

• Hybridization is the process of combining orbitals on the SAME atom.

• Adding AOs results in constructive interaction, forming bonding molecular orbitals (lower energy).

• Subtracting AOs results in deconstructive interaction, forming antibonding molecular orbitals (higher energy).

• Molecular orbital diagrams illustrate the relative energies of MOs (e.g., s-s overlap).

• Bond Order: Defined as \frac{\text{# of electrons in bonding MOs} - \text{# of electrons in antibonding MOs}}{2} .

• p orbital overlap:

  • When two p orbitals overlap along the internuclear axis (the line between the nuclei), a $\,\sigma\,$-bonding and a $\,\sigma^*\,$-antibonding orbital result.

  • Most of the electron density in a $\,\sigma\,$-bond is centered directly between the nuclei.

  • Sideways overlap of two parallel p orbitals leads to a $\,\pi\,$-bonding and a $\,\pi^*\,$-antibonding molecular orbital.

• s orbital with p orbital overlap: Gives a $\,\sigma\,$-bonding and a $\,\sigma^*\,$-antibonding MO.

• A double bond consists of one $\,\sigma\,$-bond and one $\,\pi\,$-bond.

• A triple bond consists of one $\,\sigma\,$-bond and two $\,\pi\,$-bonds.

1.15, 16 & 17 Hybridization, Drawing Three-Dimensional Molecules and Molecular Shapes & General Rule of Hybridization and Geometry

• Valence-shell electron-pair repulsion (VSEPR) theory is used to explain the molecular shape and electron-pair geometry of molecules.

• Hybridized orbitals provide a more useful method to understand bonds.

  • Hybrid orbitals result from the combination of atomic orbitals (AOs) with different quantum numbers; this process is called hybridization.

• $\text{sp}^3$ hybridization (e.g., carbon in methane, $CH_4$):

  • Mixes one $2s$ atomic orbital and three $2p$ atomic orbitals ($2s + 2p<em>x + 2p</em>y + 2p_z$) to form four $\text{sp}^3$ hybrid orbitals.

  • These four $\text{sp}^3$ hybrid orbitals are arranged to allow maximum space, resulting in a tetrahedral electron-pair geometry.

  • The ideal bond angle for $\text{sp}^3$ is $109.5^ ext{o}$.

  • Each of the four $\text{sp}^3$ hybrid orbitals can bond with a hydrogen's s-orbital, forming four C-H $\,\sigma\,$-bonds.

• $\text{sp}^2$ hybridization:

  • Three orbitals (one s and two p) combine to form three $\text{sp}^2$ hybrid orbitals.

  • The electron-pair geometry is trigonal planar (also referred to as trigonal).

  • The ideal bond angle for $\text{sp}^2$ is $120^ ext{o}$.

• $\text{sp}$ hybridization:

  • Two orbitals (one s and one p) combine to form two $\text{sp}$ hybrid orbitals.

  • The electron-pair geometry is linear.

  • The ideal bond angle for $\text{sp}$ is $180^ ext{o}$.

1.18 Bond Rotation

• Single bonds can freely rotate, giving rise to a variety of conformations (discussed further in Chapter $3$).

• Double bonds cannot rotate because rotation would break the $\,\pi\,$-bond.

• Compounds that differ in how groups are arranged around a double bond are called geometric isomers (or cis-trans isomers) and can be isolated and separated.

1.19 Isomerism

• Isomers are different compounds that share the same molecular formula.

• Constitutional (or structural) isomers differ in atom connectivity.

  • They have different names and different physical and chemical properties.

  • The number of possible isomers increases rapidly with increasing carbon number in the molecular formula.

  • There is no simple calculation to determine the number of isomers based solely on a molecular formula.

• Stereoisomers have the same molecular formula AND the same connectivity of atoms.

  • They differ only in the arrangement of atoms in space (their spatial orientation).

  • Cis and trans isomers of alkenes are specific examples of geometric stereoisomers.