Periodic table trends are helpful, but actual Electronegativity values are even more so. What about the Mg to C bond? The trend says that C is more electronegative but by how much? Knowing Mg is 1.2 and C is 2.5 (on the Pauling scale) says this is a very polar covalent bond in which the C is electron-rich and thus reactive. When we get into the weeds, e.g. B and Al in reducing reagents, knowing these numbers will explain the very different reactivities of compounds like sodium borohydride and lithium aluminum hydride.

There is a general relationship between the strength of a bond and the relative atom size and electronegativity of the elements involved. Bonds between atoms of quite different sizes tend to be weak. An example would be C-O (350 kj/mol) and C-S (260 kj/mol) single bonds. Oxygen is directly above sulfur in the periodic table so O is smaller (and closer in size to C) and more electronegative (O = 3.5 vs S = 2.5). The C-S bond is weaker and easier to break. A C-N single bond (290 kj/mol) should be favoured over a C-Br bond (275 kj/mol), which contains a larger bromine. This will be important, for example in SN2 reactions, in which nitrogen nucleophiles can displace Br on alkyl bromide substrates.

In reactions involving double bonds, sigma bonds are generally stronger than pi bonds so addition to pi bonds to produce new single bonds is usually favoured. Many addition reactions are essentially irreversible at room temperature (e.g. hydroboration), partly because of the swapping of a weaker C=C pi bond (270 kj/mol) for (cumulatively) stronger C-H (413 kj/mol) and C-B (372 kj/mol) single bonds. Expect pi bonds to be reactive in many situations.

Bonds formed between atoms that each have at least one lone pair are always going to be weak. Bonds in peroxides (RO-OR) are easily broken (O-O single bond = 146 kj/mol), which correlates with electron repulsion between lone pairs on adjacent O atoms that are close in space. Likewise, the Br-Br bond (193 kj/mol) and Cl-Cl bond (239 kj/mol) will break easily, for example in the initiation step of a radical-based process.

See here for a general discussion of bond strengths.

Acid-Base chemistry is where the discussion of mechanism begins in most Organic 1 courses. While this topic is covered in General Chemistry, its application in Organic seems to cause problems as the structures of the acids and bases involved are generally more complicated than the simple examples seen in the first year. The ideas, however, are the same.

Initially we are looking at protic acids in which the H will be bonded to a more electronegative element (e.g. OH in water and alcohols) or a bigger element (e.g. Br in HBr). HBr is a stronger acid than water because the H-Br bond is weaker and easier to break (363 kj/mol, with very different sized atoms) than the O-H bond (467 kj/mol, atoms closer in size), but also because of the relative conjugate base stabilities. Br is much better at dealing with a negative charge (extra electrons) than O because it is bigger and the charge (lone pair) has more volume to occupy. The Br anion is thus much easier to form and is much more stable than hydroxide (and later alkoxide) so this reaction favours the right-hand side completely.

Knowing actual pKa values (examples in the video below) will help in solving acid-base problems quickly. Again, trends are helpful but actual numbers will help us quantify equilibrium constants in individual steps to see if a reaction pathway (mechanism) stands a chance of actually happening. While pKa values are explained by a combination of factors (polarity, anion stability), it is convenient to have a set of numbers that include all of this background for use as a quick tool in problem solving. A more exhaustive list of pKa values will be needed later when we contemplate deprotonating weaker acids such as alkynes and amines.

Knowing what the products are in the reaction between hydroxide and HBr, we can make our first attempt at using curved arrows to show a mechanism that takes account of which bonds are formed and broken. This general idea will be used extensively throughout the Organic sequence. The electron-rich hydroxide base is attracted to the electron-poor proton on H-Br (H EN = 2.1; Br = 2.8) so they interact to begin forming a bond. Since we know only one bond needs to form and one needs to break, we can be sure this won’t be a complicated mechanism. In fact, everything can happen at once in a concerted process, which is confirmed by kinetic data, that shows the reaction to be bimolecular. Two curved arrows only will be needed here to give the reaction mechanism.

This prompts an important question; where does charge (in this case negative charge) want to be? Generally, electronegative atoms prefer holding a negative charge and larger atoms are stable with that extra lone pair. Additionally, electron delocalization (resonance) greatly stabilizes negative charges and contributes to the favourability of certain species in reactions. In acid-base situations, if the charge in one of the bases is able to delocalize, that species will generally form and be favoured. In the example below, hydroxide reacts with the carboxylic acid to give products that are heavily favoured because the negative charge in the conjugate base is able to delocalize.

More errors are made on mechanistic questions through a lack of understanding of pH than any other topic. This idea is introduced in high school and first year college chemistry classes, yet it becomes obvious at the second year level how little students actually understand about what species will, and will not, be present in a reaction flask at a certain pH level. This doesn’t need to be complicated, and we won’t be doing any calculations, but knowing what is present in base versus acid, and how it affects mechanism, will clear up a lot of confusion.

For “neutral” read inert or unreactive as there are no strong acids or bases present. We can change that by introducing an acid or a base and making the system more reactive by changing the pH. This will often be a catalytic process since the added species, acid or base, speeds up reaction but is then regenerated in the last step of the mechanism.

If we can identify the conditions in a reaction as being acidic (sulfuric acid, HBr, generic hydronium, etc.) we already know what is going to happen first; the organic substrate is going to attack the acid and the substrate will become activated. This is of major importance in acid-catalyzed reactions where neutral conditions are not sufficient for reactivity but adding an acid gets things going by activating the organic substrate. This shows up in dozens of mechanisms in Organic 1 and 2 and gives an element of predictability to how pathways will unfold. Since we are dealing with a “positive” environment, any intermediates formed during these pathways will be positively charged (oxonium ions, carbocations, etc.).

In the reaction here under acidic conditions, the substrate attacks the acid to yield a cationic oxonium ion intermediate, which may react further depending on what else is present. The environment in the next example has changed to strongly basic so the “attacker” and “attackee” change. Now the negative (electron-rich) base can attack the (electron-poor) carboxylic acid proton as a way to become stable through deprotonation. Notice how the mechanism arrows have changed in going from acidic to basic conditions. In general, organic substrates attack acids while bases attack organic substrates. This is a reliable indicator of what will happen in other mechanisms in the different pH environments.

In strong acid there cannot be any strong bases and in strong bases there cannot be any strong acids. For example, it is impossible to have hydroxide anion in sulfuric acid, and it is impossible to have hydronium in hydroxide solution (they would be neutralized immediately). In neutral conditions there are no strong acids or strong bases. Carbocations will not form under basic conditions because the base is likely to attack the substrate quickly before a carbocation can form (e.g. in the E2 process).

A note on pH -