Organic Chemistry – Functional Groups, Mechanistic Arrows & Periodic Trends

Functional Groups & Course Road-Map

• Recap of previous lecture: importance of functional groups for predicting organic reactions and properties.

• Core functional groups to be covered (alkyl halides, alcohols, carbonyls, amines, etc.) plus supporting/advanced groups.

• Goal of course: learn how to convert starting materials into value-added molecules (e.g., drug candidates) through reaction design, then verify outcomes.

Reaction Selectivity & Characterisation Example

• Case study: Taxol derivative reacted with butyl chloride + base.

  • Observed product: a single hydroxyl group replaced by a $\text{C}_4$ butyl ether.

  • Questions raised:

  • Which particular OH reacted? Why only that one? ("chemo-selectivity").

  • How do we know the reaction occurred and is clean?

• Answers come from:

  • Modern spectroscopic tools (NMR, IR, MS) to detect disappearance of starting signals / appearance of new ones.

  • Understanding intrinsic reactivity of each functional group.

Named Reaction Mentioned

• Williamson Ether Synthesis (to be covered in Week 8): nucleophilic alkoxide + alkyl halide $\rightarrow$ ether + halide ion.

Four Main Arrow Types in Organic Chemistry

• Reaction arrow ($\rightarrow$): shows conversion of reactants to products.

• Equilibrium arrow ($\rightleftharpoons$): indicates forward & reverse reactions occur (e.g. $\mathrm{HCl}_{(aq)} \rightleftharpoons \mathrm{H}^+ + \mathrm{Cl^-}$; position depends on pH).

• Resonance arrow (double-headed $\leftrightarrow$): depicts delocalisation of electrons without changing atom connectivity (e.g. benzene canonical forms).

• Curly/Mechanistic arrows (hooked, $\curvearrowright$):

  • Tail = source of an electron pair.

  • Head = destination (generally electron-poor centre).

  • Track bond-making and bond-breaking step-by-step.

Curly Arrow Example (SN2 on Bromomethane)

1. \ce{HO^-} lone pair attacks electrophilic carbon of \ce{CH3Br}.

2. Simultaneous \ce{C–Br} bond electrons depart to Br.

3. Products: \ce{CH3OH} + \ce{Br^-}.

• Demonstrates:

  • Electron-rich nucleophile seeks electron-poor site.

  • Leaving group departs with the two former bonding electrons.

  • Conservation of total lone pairs/charge across mechanism.

Electron Pushing & Pair Accounting

• Concept: visual bookkeeping of valence electrons during reactions.

• Facilitates prediction of products & design of new syntheses.

• Later lectures will formalise mechanisms for carbon chemistry.

Historical Context of the Periodic Table & Atomic Structure

• 19th-century chemists grouped elements by similar physical properties.

• 1911 Rutherford model: tiny central nucleus (protons + neutrons) surrounded by vast electron cloud. (Pea vs football-field analogy.)

• Present lecture explores how periodic trends arise from underlying electron configurations.

Periodic Table Basics

• Atomic number ($Z$) = number of protons (and, in a neutral atom, electrons).

• Atomic mass ≈ protons + neutrons; table value is weighted average over isotopes.

• Groups (vertical, 1 – 18): elements with analogous valence shells $\Rightarrow$ similar chemistry.

• Periods (horizontal, 1 – 7): progressive filling of electron shells; atomic mass generally increases left → right.

Physical Trends Across & Down the Periodic Table

Melting & Boiling Points

• Do not simply increase with atomic mass.

  • Example (Period 2): $\text{Li }(180^\circ!\mathrm{C}) \uparrow \text{Be }(1287^\circ!\mathrm{C}) \uparrow \text{B }(2450^\circ!\mathrm{C}) \uparrow \text{C }(\approx 3800^\circ!\mathrm{C})$ then sharp drop to $\text{N}_2(-210^\circ!\mathrm{C})$.

• Down a group (e.g. alkali metals) MP/BP generally decrease despite larger atoms ((\text{Li} > \text{Na} > \text{K} > \text{Rb})).

Metallic vs Non-metallic Character

• Metals: lustrous, conductive, malleable, usually solid at RT (Hg excepted).

• Metallic character increases down a group, decreases across a period.

• Periodic blocks:

  • Main-group metals (Groups 1–2 & lower part of 13–16).

  • Transition metals (d-block).

  • Metalloids (diagonal strip: B, Si, Ge, etc.).

  • Non-metals (right-hand side, esp. Groups 15–18).

Reducing vs Oxidising Power

• Strongest reducing agents: lower-left corner (e.g. \ce{Li}, \ce{Na}, \ce{K}). React violently with water (demonstrated with Na/K videos).

• Strongest oxidising agents: upper-right minus noble gases (e.g. \ce{F2}, \ce{O2}, \ce{Cl2}). Household bleach exploits \ce{Cl2} released from \ce{NaOCl}.

Atomic & Ionic Radii

Experimental Determination

• Measure internuclear distances in molecules/solids via X-ray diffraction; radius ≈ half that distance.

Neutral Atomic Radius Trends

• Decrease across a period ((\text{Li }152\,\text{pm} \rightarrow \text{F }64\,\text{pm})).

• Increase down a group due to additional electron shells.

Ionic Radii

• Cations smaller than parent atoms (e.g. $\text{Li}:152\,\text{pm} \rightarrow \text{Li}^+:78\,\text{pm}$).

• Anions larger than parent atoms (e.g. $\text{F}:64\,\text{pm} \rightarrow \text{F}^-:136\,\text{pm}$).

• Result of electron loss (greater effective nuclear pull) vs gain (added e⁻–e⁻ repulsion).

Common Ionic Charges by Group

• Group 1: $+1$

• Group 2: $+2$

• Group 13 (B, Al): $+3$

• Group 14: often $+4$ (Sn/Pb can vary)

• Group 15: $-3$

• Group 16: $-2$

• Group 17 (halogens): $-1$

Ionisation Energy ($IE$)

• $IE_1$ = minimum energy to remove the first electron (ground state).

• Periodic trend: $IE_1$ increases left → right; decreases top → bottom.

  • Alkali metals have lowest $IE_1$; noble gases highest.

• Successive $IE<em>n$ values jump sharply once a closed shell is breached (e.g. for \ce{Li}: $IE</em>1=0.52\,\text{MJ mol}^{-1}$, $IE_2=7.3\,\text{MJ mol}^{-1}$).

Electronegativity ($\chi$)

• Ability of an atom in a bond to attract shared electrons.

• Pauling scale highlights:

  • Highest: $\chi(\text{F})=4.0$ > $\chi(\text{O})\approx3.5$ > $\chi(\text{Cl})\approx3.0$.

  • Lowest: alkali metals $\chi\approx0.7!–!1.0$.

• Same qualitative trend as $IE$: up across, down the group.

• Determines bond polarity, reactivity (e.g. \ce{C–Br} polarised toward Br; carbon becomes electrophilic).

Electron Affinity (EA)

• $EA$ = enthalpy change when an atom gains an electron (negative value = energy released).

• Trends less regular than $IE$: halogens release most energy (large negative $EA$), noble gases require energy input (positive $EA$).

Key Take-Home Periodic Trends

Practical & Pedagogical Notes

• Understanding periodic trends aids prediction of:

  • Reaction mechanisms (site of attack, leaving groups).

  • Physicochemical properties (MP/BP, solubility).

  • Spectroscopic signatures (electronegativity influences NMR shifts, IR frequencies).

• Upcoming course components:

  • Week 3 onward: deeper carbon chemistry & curved-arrow mechanisms (Prof Scott Stewart).

  • Week 8: detailed Williamson Ether Synthesis.

  • Lab sessions: hands-on experiments with reactive metals (Na, K), spectroscopy, synthesis.

• Administrative reminder: labs are compulsory; email unit coordinator for timetable clashes.

These bullet-point notes encapsulate every concept, example, trend, and implication discussed in the lecture, providing a standalone study guide for upcoming assessments and practical work.