As you move beyond straight-chain hydrocarbons, you’ll notice that some carbon atoms are more “connected” than others. We classify them by substitution level:

• Primary carbon: bonded to one other carbon.
• Secondary carbon: bonded to two other carbons.
• Tertiary carbon: bonded to three other carbons.
• Quaternary carbon: bonded to four other carbons.

This classification matters because substitution affects both structure and stability, especially for charged species (see below).

Charged Hydrocarbon Ions: Structure vs. Stability

So far, every molecule we’ve seen has followed the same pattern: each carbon atom is surrounded by four electron pairs (two electrons per bond), and each hydrogen is connected by just one. That pattern is not arbitrary—we observe it because it’s based on stable molecules, i.e., those that are not reactive under normal conditions.

In many stable molecules, we often observe atoms (especially in the second period, like carbon) surrounded by eight electrons, thus filling their valence shell. (This is often called the “octet rule”.) For hydrogen, only two electrons are needed to fill its valence shell.

In methane (CH₄), this is easy to see: carbon forms four single bonds to hydrogen. Each bond shares a pair of electrons, giving carbon access to a full octet.

But what happens when a molecule gains or loses electrons? Let’s consider two modified forms of methane:

• Remove a hydrogen and both bonding electrons (i.e., an H¯), and you get CH₃⁺, a carbocation. Now carbon is only surrounded by three bonds—just six electrons, not eight.
• If you remove a proton (H⁺) from CH₄, you get CH₃¯, a carbanion. In this case, the electrons that formed the C-H bond become a lone electron pair on carbon, giving it eight electrons—but an uneven distribution of negative charge.

These are both plausible Lewis structures. But how can we tell if one is more “reasonable” than another?

This is where formal charge comes in. Formal charge helps us evaluate the electron distribution in a structure based on how many electrons an atom “owns” in the structure, compared to how many it typically brings to the table.

Formal charge assumes that electrons are shared evenly between the two atoms in the bond; for the C-H and C-C bonds in hydrocarbons, this is not a bad assumption (a point we will return to in Unit 2). So, let’s look more closely, assuming each atom in the bond “owns” one electron in that bond:

• In CH₄, carbon makes four bonds and shares one electron in each. That’s four electrons “owned” by the carbon atom, which matches the four valence electrons a carbon atom has, given its position in the periodic table. So the formal charge on carbon is 0.
• In CH₃⁺, carbon makes only three bonds—meaning it “owns” just three electrons. That’s one fewer than expected. So the formal charge on carbon is +1.
• In CH₃¯, carbon makes three bonds and has a lone electron pair, for a total of five “owned” electrons. That’s one more than expected. So the formal charge on carbon is -1.

Formal charge helps us quantify how “unbalanced” an atom is in a Lewis structure, even when the structure appears complete.

But structure alone does not tell the whole story. We also care about stability. A structure may obey the octet rule and still be relatively higher in energy—meaning it’s unstable under normal conditions.

Consider the following comparisons:

• CH₃⁺ (methyl cation) is quite unstable. The carbon lacks an octet and carries a +1 formal charge. When it reacts with H¯ to form CH₄ (methane), about 1340 kJ/mol of energy would be released
  • i.e., the CH₃⁺ + H¯ ⟶ CH₄ reaction is energetically favorable by ~1340 kJ/mol.
• CH₃¯ (methyl anion) has a full octet, but it concentrates negative charge on a small atom with no way to stabilize it. When it reacts with H⁺ to form CH₄ (methane), an even greater 1780 kJ/mol of energy would be released
  • i.e., the CH₃¯ + H⁺ ⟶ CH₄ reaction is energetically favorable by ~1780 kJ/mol.
• In the tert-butyl carbocation, the positively charged carbon is surrounded by three methyl groups (it is bonded to three other carbon atoms). When it reacts with H¯ to form isobutane, about 1110 kJ/mol of energy would be released
  • i.e., the C(CH₃)₃⁺ + H¯ ⟶ HC(CH₃)₃ reaction is energetically favorable by ~1110 kJ/mol.
  • The tert-butyl carbocation is still unstable, but it is less so compared to the methyl cation because the electron density of the surrounding atoms (and thus bonds) help to stabilize the positive charge.

In short: formal charge tells us where charge is (usually, but not always because the assumption that the electrons are shared evenly between the two atoms in a bond is not always valid), while stability depends on how well that charge is handled.