But, How Do I Know The Structure (and why should I care)?

Why should I even read this?

To be blunt, it is not possible for you to pass this course, or probably even the first exam, if you can’t draw the structure of a compound, given the formula and/or name for that compound. Want to see what I mean? Draw the structure of the compound of condensed formula CH3COCH(NH3)CH2CHPhCO2CH(CH3)2. Got that one? Now try acetonitrile. In Organic Chemistry, you will learn, and demonstrate that you have learned, many things about the structures, properties, and reaction patterns of about 11 different groups of compounds. It isn’t possible to do this if you can’t tell which group the compound you’re given fits into. Few Organic textbooks do a very good job of reviewing how to draw structures from formulae, and you most likely haven’t had to do this since about mid-first semester in General Chemistry.

Where To Start

Drawing the structure of something, given a condensed formula, mostly goes back to the use of common sense and the periodic table. Just because this is a ‘new’ chemistry class doesn’t mean you throw out everything you learned in earlier chemistry classes! You will find yourself constantly challenged to remember and use what you learned (or were supposed to learn) in GenChem. Since most students really don’t retain a whole lot of their GenChem this is a problem! So let’s go back to the Periodic Table for a little review of bonding patterns. (You know, bonding? The term that describes what holds atoms together to make molecules and compounds?)

You must, emphasize must, know the first three rows of the Periodic Table by heart. You may not have been required to learn it before, but there is no way to succeed in Organic unless you know it now. You will not, not, not get a Periodic Table on exams. You need to know the location of certain elements and you must be able to use both element symbol and element name interchangeably. Remember, the Periodic Table is set up as it is because, in general, every atom within a given horizontal row has certain things in common, and every atom within a given vertical column has certain things in common. A compact version of the Periodic Table for most of the atoms you need to know is shown below.


Abbreviated Periodic Table of the Elements




















































These concepts and terms are discussed in another handout, but to briefly review, the horizontal rows of the Periodic Table are called "Periods." Each Period represents a different quantum energy level. Within each Period there may be one, two, three or four sublevels, which are designated s, p, d, and f, each representing a different, increasing, energy level for the atom’s electrons. All of the elements/atoms in a given row are considered to have sets of electrons, on a relatively atomic large scale, at the same energy level, the principle quantum level (1, 2, 3 etc.). Each numerically higher quantum level is built on top of the one before, so there is an increasing number of electrons at different quantum energy levels. As a general rule, only the outermost, or valence, electrons are shared or transferred to result in a bond with another atom.

The vertical columns of the Periodic Table are called the "Groups." Some Tables use Roman Numerals (I, II, etc.) while others use Arabic (also called Hindu) numerals (1, 2, etc.). It is sometimes easier to keep Groups and Periods clear when you use different numerals. Groups are organized (generally) such that elements with similar properties are in the same group. You need to know the common names for some of the groups: Group IA are alkali metals, IIA are alkaline earth metals, VIIA are halogens, IB, IIB, VIB and VIII are transition metals.

Bonding Patterns: To Octet or not Octet

Bonds may be ionic or covalent. An ionic bond arises when one atom fully transfers a valence electron to the other atom, and the two atoms are held together by the electrostatic attraction of opposite ionic charges. There is no electron sharing in a fully ionic bond. A covalent bond is exactly what the word says, co (both) valent (valence electron), or, each atom contributes one valence electron to the bond and both electrons are then shared to some degree between the two atoms. Covalent bonds range from being an absolutely equal sharing (fully covalent), to being almost fully ionic, or where the sharing is really unequal, but not so far as an electron transfer. Covalent bonds with unequal sharing are called polar covalent bonds. Nearly all the bonds in organic compounds are covalent, and we will focus on those bonds here.

You probably learned something called the Octet Rule, a simplification of the concept that atoms will react (gain, lose, or share electrons) in order to have the same number of valence electrons as the nearest noble gas element. The Octet Rule says the sum of all the shared and unshared valence electrons about an atom must total 8 to have a stable species, which translates to 4 covalent bonds. Sorry, it doesn’t work for all atoms. Hydrogen follows the Duet Rule. Boron and Aluminum follow the Sextet Rule. Sulfur is confusing, as it can fit into the Octet, Dodectet (12) or Octadectet (18) Rule!

What to do? Use common sense and what you’ve already learned to predict correct bonding patterns. Quick, what’s the formula for water? Just about everyone answers "H2O," or 2 Hs covalently bonded to O. So you remember oxygen forms two covalent bonds when neutral. Can you do ammonia? "NH3," or 3 Hs covalently bonded to N, or 3 covalent bonds to N when neutral. See a pattern? Now predict carbon.

There are some generalities to seeing when the Octet Rule does not apply. Clearly, Groups IA, IIA, and IIIA cannot accommodate 8 electrons, as they don’t have enough valence electrons to form that many bonds. Molecules with an odd number of total valence electrons also don’t follow the octet rule, such as NO. Finally, and most annoying to Organic Chemistry students, are those elements which have an expanded valence shell. Period 3 and higher elements can use the energetically low-lying 3d sublevel (atomic orbitals) to increase the number of bonds formed. These elements will have more than an electron octet. The compounds you will encounter the most often will be oxyacids, i.e. H2SO4 (sulfuric acid), and reagents like PCl5.

Keep in mind that we are mostly concerned with drawing structures for compounds made up of Period 2 atoms, and primarily those containing carbon. Knowing the Group number for these atoms gives you their usual covalent bonding patterns. For Groups IA, IIA and IIIA the number of bonds that can be formed using the valence electrons of those atoms is equal to the group number. For carbon (C, Group IVA) the maximum number of bonds is the magic 4, or electron octet. Now you have to consider the period as well as the group number. For neutral Groups VIA – VIIA (N, O, F with no formal charge) the number of bonds is equal to 8 — Group #.

Group properties also help to predict covalent versus ionic bonding. The farther apart the two bonded Groups are, the more likely the bond will be ionic, and the closer the two Groups are, the more likely the bond will be covalent. A rule-of-thumb is that if the two elements are more than 4 groups apart, the bond will be ionic. This all becomes clear when you look at differences in electronegativity (the ability of an atom to attract electrons to itself). The greater the difference in electronegativity (D EN) the more the electrons in the bond are pulled towards the most electronegative atom. At some point the difference is so great that the electrons are actually transferred rather than shared. When D EN is equal to or more than 1.7, the bond is ionic. Likewise, you can predict that the further apart the row, the less likely the elements are to covalently bond to form a stable compound. Boron trifluoride is stable while boron triiodide is not. Hydrogen is technically located in Group IA as it has only one valence electron. In terms of all of its periodic properties, though, it should be in Group IVA. Hydrogen chloride (H—Cl) and lithium hydride (Li—H) are covalently bonded. Hydrogen is more like carbon than like lithium in many bonding patterns, and there are two, not one, reactive hydrogen ions, H+ and HG .

Charges and Charges

There are two kinds of charges to watch out for when drawing structures. (No, not MasterCard and Visa.) Ionic charges are formed when an element loses one (or more) electrons to form a cation, or when an element gains one or more electrons to form an anion. You will mostly encounter Group IA and IIA cations i.e., Na+, K+, Mg2+, and Ca2+, and Group VA, VIA and VIIA anions, i.e., Cl, Br, HO, and H2N. The other ionic species you will use a lot are carbocations, and carbanions, protons (H+) and hydrides (H). Ionic charges are a form of electron "bookkeeping" when electrons are actually lost or gained by an atom.



Formal charges are a way to do "electron bookkeeping" for covalent bonds. These do not indicate any "real" ionic charge. There may or may not be an overall ionic charge when there are formal charges within a molecule. The sum of the formal charges within a molecule must equal the total ionic charge. Assigning a formal positive charge to an atom does not mean that atom actually is positive. The rule for writing formal charges is simple. There should be as few formal charges as possible, and when formal charges must be assigned, they should be as low as possible. The distribution of formal charges is such that the most electronegative atom gets the negative charge(s). Formal charge (FC) = Group # — (# lone pair electrons + ½ # bonding electrons).

Beyond Lewis Structures

Instead of always counting electrons and struggling with Lewis structures, there is a fairly quick way to remember bonding patterns. Look at the diagram below.







Li Be B C N O F

# of bonds (neutral) 1 2 3 4 3 2 1

# bonds when +1 FC 0 3 4 3 2

# bonds when –1 FC 4 3 2 1 0


If you always relate back to carbon with 4 bonds when neutral and 3 when charged, you can use the idea of a pyramid to predict the rest of the atoms. This holds true for trends in electronegativity, too.

Drawing Structures

You’re finally ready to draw some structures. Organic structures are pretty easy because there are some other generalizations you can use too. True Lewis Structures show bonds as pairs of dots. Kekulé Structures show bonds as lines and use dots to represent non-bonding electrons. We use Kekulé structures but call them Lewis.

  1. Bonds are always made up of pairs of electrons, not single electrons.
  2. Initially figure there are at most 4 ‘directions’ of electrons around each atom except H. It is often easiest if you draw out each atom showing the electrons in up to but not more than four directions.
  3. Paired up electrons around an atom (drawn before connecting to another atom) are less likely to be part of a covalent bond than an unpaired electron.
  4. The maximum number of bonds equals (total of all valence electrons) / 2. The actual number of bonds is most likely to equal (total of all unshared electrons) / 2. If the total of unshared is odd, go to the lower number (i.e., 17 drops to 16) and look for a formal charge.
  5. Usually the way you see the compound written out is the way the bonding goes for all atoms except H.
  6. Subscripts mean that number of the atoms or groups preceding the subscript are present and are bonded to the atom either just preceding or following the subscripted atom/group.
  7. Parenthesis ( ) are used to clarify a particular bonding or to keep several atoms together as a group.
  8. If there is more than one C, then usually the C are bonded to each other.
  9. Draw all single bonds before you go looking for double or triple bonds or cyclic structures.
  10. There may be more than one correct Lewis Structure for a molecule.
  11. Above all, use common sense and remember all the compounds you’ve drawn so many times before!




total of unshared e = 8: 8/2 = 4 total of unshared e = 8: 8/2 = 4

so look for 4 bonds to form so look for 4 bonds to form



Note the 3 middle C don’t have ‘4’ things on them, so

Look for multiple bonds


Total of unshared e = 16

16 / 2 = 8 bonds


31 unshared e drops to 30, 30/2 = 15 bonds & a –1 FC