Organic Chem Basics

Preparing for Organic Chemistry

Being ready for Organic 1 is key.

General Chemistry Concepts…

You need to have a good understanding of why the periodic table is organized the way it is. Knowing roughly where the main atoms reside in the table will be a huge help. 

Knowing why the elements are organized in rows and groups based on their electronic structures will allow the student to make predictions on properties when the elements show up in Organic molecules.

To get started: Know the important elements early on:

  • You don’t need to know all of the elements; to get started you should focus on the top of the table and elements that actually get used early on in Organic 1. Elements (e.g. transition metals) will be added later as they are needed for new reactions.
  • Focus on Carbon being in the middle of this abbreviated table; this is important as C has a “middle” electronegativity (2.5). Elements to the left of C are less electronegative; those to the right are more electronegative. This will dictate types of bonds formed.
  • Metals on the left give away electron(s) to form ionic bonds, those in the middle share, those on the right can accept electrons or share to form ionic or covalent bonds. You learned this in General Chemistry and it still applies in Organic Chemistry.
  • Before you begin Organic 1 you should know where these elements are on the table and why they are organized this way. This will save time when making decisions on how the elements will interact based on their electronegativities and whether they can share of transfer electrons. If you are looking at a periodic table during an exam to find out where an element is, you are wasting time that should be spent elsewhere.
  • Notice that some elements are crossed out. We don’t talk about Berylium (it’s rare) and we don’t make molecules from the inert (Noble) gasses (because they are inert with a full octet). The rest of the elements in the above table will do what they did in General Chemistry so you should be on top of that as you prepare for Organic 1.

The attraction of an atom or group for electrons. Knowing some actual numbers will be more useful than just an appreciation of trends. This will help you to decide which type of bonding is occurring in molecules.

Know about Electronegativity

  • This idea, which you first studied in General Chemistry, is so important that it will help to know some actual numbers. Instead of just learning a general trend, by being more specific you will be able to solve problems more confidently.
  • Electronegativity (EN) of elements increases from left to right in the Periodic Table in ~0.5 increments heading to F, which has the highest electronegativity value of 4.0 (on the 4.0 scale). Going from left to right, valence electrons are added to the same energy level while the nucleus gains one extra proton per group. The more positive nuclei will attract the valence electrons more tightly, which affects atomic size and then reactivity. Going down the Table the valence electrons get further away as electrons go into higher orbitals (3s, 4s, etc.) so the influence of the nucleus fades.

Pauling Scale electronegativity values for the important early elements.

  • Knowing the numbers means you are able to decide what kind of bonding will occur in between different atoms. Very different EN values, e.g. Na @ 0.9 and Cl @ 3.0 are only ever forming ionic bonds. C @ 2.5 and H @ 2.1 will only form covalent bonds with each other through sharing of valence electrons. There is no absolute cutoff difference, however.
  • Once molecules are formed, you will then be able to decide which areas within are non-polar or polar, which is essential for understanding reactivity patterns. Anions will be electron-rich and go after electron-poor atoms in polar-covalent areas.

The number of bonds an atom may form, and the number of lone pairs that are (or are not) present, will dictate the reactivity of organic molecules and their ability to act as nucleophiles or electrophiles. 

Understand Valence Bond Possibilities

  • This comes straight from the Periodic Table and having a working knowledge of it is essential. An atom with one electron in its valence shell will form one bond, most likely by giving that electron away (think Li, Na, K). Atoms with seven electrons will form one bond to gain the octet (think F, Cl, Br). An atom with four valence electrons (carbon!) will need to form 4 bonds maximum to achieve the octet.
  • Electronegativity (EN) of elements dictates what type(s) of bonds are formed but the valence atomic structure governs how many. The table below gives an idea of what to expect. Just remember that 8 is the desired number for most of the upper key elements. Smaller atoms (H, Li. etc.) don’t need 8 as 1s shell can be filled to give a full valence.

Number of bonds formed by each of the important early elements.

  • Carbon is very flexible, due to its middling EN. It forms 4 bonds in its stable (neutral) molecules but will be able to cope with three bonds temporarily in intermediates. C will never form 5 bonds in Organic 1 or 2. Boron (B) and Aluminium (Al) are unique at this stage in that they form 3 bonds in neutral molecules but will form 4 (for the octet) during the various reactions in which they feature. They are used as Lewis Acids.
  • Atoms to the right of C will have lone pair(s) in their neutral molecules and will be able to share with other atoms to form common species such as hydronium or ammonium in which the central O or N atom has a positive charge. Same as in General Chemistry.

Related to periodic trends, the size of atoms will play a very important role in their 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 are required to put together all of the concepts studied thus far, including the mechanism of the proton transfer process.

…are also Organic Concepts

Just because the name of the course changes the basics do not; same if you move on to Biochemistry and/or Physical Chemistry later. Don’t be coming in to Organic cold; give yourself a chance to build on what you already know.

Almost all of the elements used early in the undergraduate Organic Chemistry sequence are at the top of the periodic table and thus are in search of the perfect electronic “octet” in their valence shell.

  • The Periodic Table is an organization of elements in their atomic form before they react to form molecules. Apart from the Nobel (inert) gasses, which already have the complete valence shell, elements lose or share electrons to match the Noble gasses.
  • On the left of the table (think Li, Na, Mg), atoms have 1 or 2 valence electrons (and low EN values) so it will be easier to lose electrons than share or pick them up. Atoms on the right (think O, F, Br) will pick up electrons to achieve the octet. Those in the middle will share electrons in covalent bonds. How many electrons each atom needs depends on how many they have to begin with, which comes from the Periodic Table.
  • Don’t forget about non-bonding electrons (lone pairs, l.p.), which count towards the octet: N (also P) has 1; O (also S) has 2, F (also Cl, Br, I) has 3.

How each element will get to the nearest Noble gas octet.

  • How an atom gets to the octet is a function of its own electronegativity and what element(s) it is bonding to. Each of the elements above is capable of forming either covalent or ionic bonds depending upon the other elements involved.
  • Do not worry about elements that can expand their octet (S, P, etc.) at this point; that will come later in the Organic courses. For now the abbreviated table above shows you how each of the important elements will achieve the octet.

Bonds are a result of the relative electronegativities of the atoms involved; very different electronegativities (e.g. Na and Cl) means ionic bonds through electron transfer, while close electronegativities means sharing in covalent bonds.

  • In General Chemistry you spent a lot of time studying ionic salts, their composition, and their stoichiometry in chemical reactions. While that material is useful it can also be confusing. In the Organic classes we start simple but we need to expand on your understanding of which bonds are most likely to be formed in various situations. Most of this relies on understanding relative electronegativity (EN) values (see earlier).
  • In the Organic courses we expand upon the basic idea of “ionic” and “covalent” bonds and invoke the “polar covalent bond” as being essential for reactivity. Ionic bonds will be obvious as they are formed between atoms on the left and right of the Table. When atoms are closer together (and thus have similar EN values) they share bond electrons in covalent arrangements. Some difference in EN values then leads to polar covalent.

The main bonding patterns seen in Organic Chemistry.

  • The question will be, “where are the electrons” in each bond? With large differences in EN the electron(s) will be associated completely with the more EN atom. For moderate differences the electrons are shared but closer to the more EN atom to give a dipole. For small EN differences the electrons are equally shared and no dipole exists.

Molecule shapes will be dictated by the number of sigma bonds and lone pairs involved; 4 sigma bonds = tetrahedral (with slight variations if lone pairs replace sigma bonds); 3 sigma bonds = trigonal planar (i.e. flat), and 2 pairs = linear. 

  • The shapes of molecules depend upon the number of valence pairs that are present in sigma bonds and lone pairs. Shape has nothing to do with pi bonds. Sigma bonds and lone pairs will arrange themselves to be as far apart as possible from each other (since negative electrons repel). This gives tetrahedral, trigonal planar, and linear shapes.
  • Starting with Carbon, the molecule methane (C bonded to 4 hydrogens) has only sigma bonds, which repel each other. The tetrahedral shape is simply the furthest away these four bond pairs can get from each other. When pi bonds are introduced it is still the single bonds that dictate shape; trigonal (3) planar for 3 sigma and linear for 2.
  • Moving to Nitrogen and Oxygen we recognize the introduction of lone pair(s) which are known to take up more volume than bond pairs and so change these shapes slightly. For ammonia (N bonded to 3 hydrogens) and water (O bonded to 2 hydrogens) the tetrahedral shape still applies but with slightly different bond angles.
  • Similarly, for N and O in double bonds the shape is roughly trigonal planar but the associated lone pairs distort the angles away from 120. The molecules will, however, still be flat overall. For the linear molecules the N equivalent (nitriles) has the lone pair opposite to the alkyl group attached with the molecule still being linear overall.

This is mostly governed by bond strengths and atoms in molecules being electron-rich or electron-poor; electron-rich atoms (i.e. bases or nucleophiles) attack electron-poor atoms (i.e. Lewis acids or electrophiles). Atoms and molecules want to become more stable and have a variety of pathways (mechanisms) to low energy states; some pathways are reversible, which leads to equilibria and the application of Le Chatelier’s principle. Kinetics relates to how fast a reaction occurs and how high activation barriers will be; simple ideas like a crowded environment means slower reactions while accessible environments (i.e. easier to get to) means faster reactions and lower Eact.

  • As in General Chemistry, molecules are going to react with other molecules so that the whole system becomes more stable. Weak bonds get swapped for stronger bonds while charge (positive or negative) finds its way to the best location possible. These ideas are used throughout the Organic sequence to explain how chemicals change.
  • What constitutes a reactive molecule? Weak bonds are a good place to start. Numerous reactions feature reagents such as Br-Br, HO-OH and similar which have atoms with multiple lone pairs directly next to each other. These bonds will always be broken along the way during reactions to give more stable outcomes. Likewise, when atoms of quite different size bond together those bonds tend to be weak and easily broken.
  • When an atom is missing electrons from the octet (e.g. B, Al, C+) they will be reactive towards electron-rich species. Positively charged carbon (i.e. carbocations) will feature regularly in Organic 1 and 2 and serve as electron-poor reactive intermediates.
  • Their reactions as Lewis acids with electron-rich Lewis bases gives molecules with full octets. Also, when negative charge may be swapped from less stable (maybe smaller) atoms to one able to stabilize the charge better, this will result in predictable reactions.

Take our course on General Chemistry Basics needed for Organic 1

Familiar Acid-Base Chemistry:

Consider the simple acid-base reaction between HCl and NaOH shown in Equation (i) that is studied extensively in General Chemistry. We know that the products are NaCl (salt) and water, that heat is given off so the reaction is known to be exothermic, meaning the products are more stable than the reactants. Everyone beginning Organic Chemistry has done that reaction at least once in General  Chemistry lab. We can use this reaction as a starting point for Organic reactions and mechanisms.

In Organic Chemistry we need to dig deeper and ask why this reaction occurs, why it is exothermic, and how might it be described in terms of the required bonds being formed and broken. Most of the answers will employ the simple concepts listed above.

We know that NaOH and NaCl contain ionic bonds because of the big differences in electronegativity between Na (0.93) and O (3.5) an Cl (3.0). This means Na metal gave away its single valence electron to become Na+ in each of those salts with O and Cl being able to complete their octet by picking up that electron. We notice that Na+ on the left of Equation (ii) is still Na+ on the right so it has not changed chemically and is therefore ignored as a spectator ion in the eventual mechanism.

So why is the right-hand side of this equation favoured so heavily and how might we describe the bond-forming and breaking events that need to occur to turn reactants into products? From first principles we can talk about reactivity and stability of the different species using bond energies and considering where the electron excess (the negative charge) will be more stable. In Equation (iii) the water product has a stronger covalent bond than in HCl and the negative charge prefers to be on the large chloride ion.

The products are more stable than the reactants, which explains why heat is released during this reaction, but what causes the starting materials to react the way they do? Why do O and H bond to give water and Na+ swaps its anion to become NaCl? The why and the how will lead us to a reaction pathway (a mechanism!). Again, electronegativity will explain most of this. The O in hydroxide is negatively charged and electron-rich while the H in HCl has a slight positive charge and is electron-poor as seen in Equation (iv).

We find that the electron-rich part of NaOH (the O) attacks the electron-poor part of HCl (the H) with the O acting as electron-donor (base) and the H acting as electron-acceptor (acid). The extra lone pair from O becomes the new bond pair in water and the covalent bond between H and Cl breaks, to deposit a fourth lone pair onto chloride, thus avoiding breaking the octet rule. The mechanism for this process may be described using the Organic Chemistry “curved arrow” notation as described in Equation (iv).

Notice what we did here; we identified the bonds that needed to be formed and broken and then considered which of the involved atoms would be electron-rich and which would be electron-poor. The electron-rich part of the NaOH molecule (O, the basic center, or later, the nucleophile) then attacked the electron-poor part of the HCl molecule (H, the acid, or later, the electrophile). This type of analysis will get you far in Organic Chemistry even when the molecules and mechanisms get to be quite complicated.

Expanding to Organic Acids & Bases:

Consider the reaction in Equation (v) below. We have to decide which starting material is the acid, which is the base, which bonds need to form and break, which side of the reaction is preferred, and which mechanism arrows need to be applied. That sounds like a lot to remember but it’s a process that becomes easier with more practice. To begin, the molecule with the metal involved is usually the base since the atom next to the metal is negative and thereby electron-rich. Here the alcohol is the acid and sodium amide (NaNH2) the base.

Since the Na+ does not change we can ignore it again and focus on the O-H bond that needs to break and the N-H bond that needs to form. Knowing that the O is better with the negative charge, since O it is more electronegative than N, the right side will be favoured. We correlate this idea with pKa values early in the first semester. Now the mechanism may be written as in Figure (vi) with the electron-rich N attacking the electron-poor H on O and giving the observed products.

The discussion of acid-base chemistry from the Organic Chemistry perspective usually comes early in the undergraduate course sequence and provides a platform for the investigation of some 120 reactions and mechanisms over the two semesters of study. Use the GenChemBasics PDF file below to review the main topics before beginning Organic 1.

GenChemBasics