Knowing why the elements are organized in rows and groups based on their electronic structures lets you predict properties when those elements show up in Organic molecules.

• You don’t need every element; focus on the top of the table and the ones used early in Organic 1. Transition metals get added later as new reactions call for them.
• Carbon sits in the middle of this abbreviated table with a "middle" electronegativity (2.5). Elements to its left are less electronegative; those to the right are more so; this dictates the bonds formed.
• Metals on the left give away electrons to form ionic bonds, those in the middle share, and those on the right accept or share electrons to form ionic or covalent bonds.

Before you begin Organic 1 you should know where these elements sit on the table and why; it saves time when deciding how elements interact based on electronegativity and whether they share or transfer electrons. Hunting for an element mid-exam wastes time better spent elsewhere.

Notice a couple of elements are crossed out: we don’t talk about Beryllium (it’s rare) and we don’t build molecules from the inert Noble gases (they’re already full-octet). Everything else behaves as it did in General Chemistry.

The attraction of an atom or group for electrons. Knowing actual numbers; not just the general trend; helps you decide which type of bonding is occurring in a molecule.

• Electronegativity (EN) increases left to right in ~0.5 increments, topping out at Fluorine (4.0 on the Pauling scale). Valence electrons join the same energy level while the nucleus gains a proton per group, pulling them in tighter.
• Going down the table, valence electrons sit in higher orbitals and further from the nucleus, so its influence fades.

• Very different EN values (Na @ 0.9, Cl @ 3.0) only ever form ionic bonds. Similar values (C @ 2.5, H @ 2.1) only form covalent bonds through sharing. There’s no hard cutoff.
• Once a molecule is formed you can call regions non-polar or polar, essential for reactivity; anions are electron-rich and go after electron-poor atoms in polar-covalent regions.

The number of bonds an atom may form, and the lone pairs present or absent, dictates the reactivity of organic molecules and their ability to act as nucleophiles or electrophiles.

• An atom with one valence electron forms one bond, usually by giving it away (Li, Na, K). Seven valence electrons also means one bond, to complete the octet (F, Cl, Br). Four valence electrons; carbon; needs up to 4 bonds for the octet.
• EN dictates the type of bond; valence structure dictates how many. Smaller atoms (H, Li) don’t need 8 since the 1s shell fills with 2.

• Carbon’s middling EN makes it flexible: 4 bonds when stable, 3 temporarily in intermediates, never 5 in Organic 1 or 2. Boron and Aluminium form 3 bonds neutrally but 4 during reactions; useful as Lewis acids.
• Atoms right of carbon carry lone pairs in their neutral form and can share them to form species like hydronium or ammonium, where the central atom carries a positive charge.

Related to periodic trends, atomic size plays a key role in an atom’s ability to hold a negative charge; as conjugate bases and as leaving groups in substitution and elimination reactions.

First studied in General Chemistry, Organic acid-base reactions are the first time students put together every concept studied so far, including the mechanism of the proton transfer process.

Just because the course name changes, the basics don’t; same goes if you move on to Biochemistry and/or Physical Chemistry later. Don’t come into Organic cold; build on what you already know.

Almost every element used early in the undergraduate Organic sequence sits near the top of the periodic table, in search of the perfect electronic "octet" in its valence shell.

• The periodic table organizes elements in their atomic form before they react. Apart from the Noble gases, elements lose or share electrons to match them.
• Left-side atoms (Li, Na, Mg) have 1–2 valence electrons and low EN, so it’s easier to lose them than share or pick up more. Right-side atoms (O, F, Br) pick up electrons. Middle atoms share in covalent bonds.
• Don’t forget non-bonding lone pairs, which count toward the octet: N (and P) has 1; O (and S) has 2; F (and Cl, Br, I) has 3.

Bonds result from the relative electronegativities of the atoms involved. Very different EN values (e.g. Na and Cl) mean ionic bonds via electron transfer; close EN values mean sharing in covalent bonds.

• In Organic we expand "ionic" and "covalent" to include the "polar covalent" bond; essential for reactivity. Ionic bonds form between atoms on opposite sides of the table; similar-EN atoms share electrons covalently, with some difference giving polar covalent.

• Where are the electrons in a given bond? Large EN differences mean the electrons sit fully with the more electronegative atom. Moderate differences share them closer to that atom, giving a dipole. Small differences share them equally; no dipole.

Molecule shapes are dictated by the number of sigma bonds and lone pairs: 4 sigma bonds give tetrahedral (with variations for lone pairs), 3 give trigonal planar, and 2 give linear.

• Sigma bonds and lone pairs repel and arrange to be as far apart as possible. Methane’s four sigma bonds give the tetrahedral shape; introducing pi bonds doesn’t change this; sigma bonds still dictate shape.

• Lone pairs (as in ammonia and water) take up more volume than bond pairs, distorting the tetrahedral angles slightly. Same for trigonal-planar N/O double bonds and linear nitriles; the molecule stays flat or linear overall.

Reactivity is mostly governed by bond strengths and whether atoms are electron-rich or electron-poor; electron-rich bases/nucleophiles attack electron-poor Lewis acids/electrophiles, moving toward more stable, lower-energy states.

• Weak bonds are the first thing to break. Reagents like Br-Br and HO-OH pair atoms with multiple adjacent lone pairs; always broken along the way to a more stable outcome. Bonds between very differently sized atoms are also typically weak.

• Atoms missing octet electrons (B, Al, C+) react readily with electron-rich species. Positively charged carbon; carbocations; are common electron-poor reactive intermediates throughout Organic 1 and 2.

The products are NaCl (salt) and water; heat is given off, so the reaction is exothermic; the products are more stable than the reactants. Everyone starting Organic has run this reaction at least once in General Chemistry lab.

In Organic Chemistry we dig deeper and ask why this reaction occurs, why it’s exothermic, and how to describe it in terms of bonds formed and broken.

NaOH and NaCl contain ionic bonds because of the large EN difference between Na (0.93) and O (3.5)/Cl (3.0). Na gives away its single valence electron to become Na+, letting O and Cl complete their octets.

Na+ on the left is still Na+ on the right; chemically unchanged, so it’s ignored as a spectator ion in the eventual mechanism.

Why is the right-hand side favoured, and how do we describe the bond-forming and breaking events? Using bond energies, we consider where the excess negative charge is more stable.

The water product has a stronger covalent bond than HCl, and the negative charge prefers to sit on the larger chloride ion; explaining why heat is released.

The O in hydroxide is negatively charged and electron-rich; the H in HCl carries a slight positive charge and is electron-poor.

The electron-rich O attacks the electron-poor H. O’s extra lone pair becomes the new bond in water; the H-Cl bond breaks, depositing a fourth lone pair onto chloride and preserving the octet rule.

We identified the bonds that needed to form and break, then decided which involved atoms were electron-rich and which were electron-poor.

This type of analysis will get you far in Organic Chemistry, even as the molecules and mechanisms get more complicated.

We have to decide which starting material is the acid, which is the base, which bonds form and break, which side is preferred, and which mechanism arrows apply. It gets easier with practice; the molecule with the metal is usually the base, since the atom next to it is negative and electron-rich. Here the alcohol is the acid and sodium amide (NaNH₂) the base.

Na+ doesn’t change, so we focus on the O-H bond that breaks and the N-H bond that forms. O holds the negative charge better than N (more electronegative), so the right side is favoured; the same idea behind early pKa comparisons.

Organic acid-base chemistry usually comes early in the undergraduate sequence and sets up the investigation of roughly 120 reactions and mechanisms across two semesters.