Formal Charge, Resonance & Organic Nomenclature

Formal Charge, Resonance & Organic Nomenclature

Lecture by Dr Jodie Johnston (CHEM 112 – Ch 5 & Ch 2)

E-mail: jodie.johnston@canterbury.ac.nz

1. What are Resonance and Formal Charge?

Resonance and formal charge are like two tools that help us understand how electrons are spread out in molecules. Lewis structures (our usual way of drawing molecules) sometimes can't show the full picture of where electrons really are, especially when they're not stuck between just two atoms.

• Formal charge helps us count electrons around each atom in a molecule. It's important for figuring out which way of drawing a molecule (called a 'resonance contributor') makes the most sense and is the most stable. We try to have as few charges as possible, and if there are negative charges, they should be on the atoms that like electrons the most (electronegative atoms).

• Resonance means that for many molecules, the electrons aren't just in one place. Instead, they are 'shared' or 'spread out' over several atoms. We show this by drawing a few different possible Lewis structures (called 'resonance contributors'). The real molecule is actually a mix or an 'average' of all these contributors, called a resonance hybrid. This spreading out of electrons makes the molecule more stable and gives it special properties.

In CHEM 112, we use Lewis structures, but later courses (CHEM 211/212) will teach you a more advanced way of looking at electron distribution using 'molecular orbitals'.

2. Why Do We Need Resonance? (Clues from Experiments)

1. Bonds come in different types: single, double, and triple. They have different lengths (e.g., C-C is longer than C=C, which is longer than C≡C) and different strengths (triple bonds are strongest, single bonds are weakest). These differences are well-known.

2. Sometimes, when we test molecules in the lab (using methods like X-ray crystallography for bond lengths or spectroscopy for bond strengths), the results don't match what we expect from a simple single or double bond. Instead, the bonds show intermediate lengths or strengths. For example, if a molecule has a bond that looks like a single bond in one part and a double bond in another, but experiments show they are both the same length somewhere in between, then resonance is happening.

3. A famous example is benzene. All the carbon-carbon bonds in benzene are the same length, about $1.39 \text{ \AA}$ (angstroms). This length is exactly halfway between a typical C-C single bond ($1.54 \text{ \AA}$) and a C=C double bond ($1.34 \text{ \AA}$). This tells us that the electrons are spread out evenly over the whole ring, not in specific single or double bonds.

3. A Quick Look at Molecular Orbitals (Very Basic)

• A sigma (σ) bond forms when atomic orbitals overlap head-on. The electron density is mostly between the two atoms.

• A pi (π) bond forms when unhybridised p-orbitals overlap side-by-side. The electron density is located above and below the main line connecting the atoms.

• When several atoms next to each other have these p-orbitals lined up, the pi electrons aren't stuck between just two atoms. They can spread out (delocalise) over many atoms. This creates 'molecular orbitals' that cover a larger part of the molecule.

• When electrons can spread out like this, the molecule becomes more stable (needs less energy). This increased stability is the main reason we see the resonance effect in many molecules.

4. Resonance Contributors vs. The Real Molecule (Hybrid)

• All resonance contributors have the same atom connections. They are not different molecules that can turn into each other (those are called isomers). Only the electrons (specifically pi electrons and lone pairs) move around between contributors.

• Curved arrows always show electron movement, never atom movement. The atoms stay in their exact positions.

• The resonance hybrid is the true structure. It's an average of all the resonance contributors, and it doesn't change back and forth. The most stable contributors contribute more to this average.

We decide which contributors are more stable by looking at:

- **Fewer charges:** Structures with fewer positive and negative charges separated from each other are more stable.

- **Full electron shells (octets):** Structures where all atoms (especially carbon, nitrogen, oxygen, and fluorine) have a full eight electrons around them are more stable. Carbon should *always* have a full octet in important contributors.

- **Negative charge on electronegative atoms:** If there are charges, negative charges are better placed on atoms that like electrons more (like oxygen instead of carbon). Positive charges are better on atoms that are less electronegative.

- **More bonds:** Structures with more covalent bonds are generally more stable.

• The contributors themselves don't actually change into each other over time. The molecule only exists as the hybrid. The double-headed arrow ($\leftrightarrow$) between resonance contributors simply means "is represented by," not "is in quick back-and-forth".

5. Benzene & Aromaticity

Benzene is a special type of compound called an 'aromatic' compound. It's super stable and behaves in unique ways.

• We can draw benzene using two 'Kekulé' structures, which show alternating single and double bonds. These are the main resonance contributors.

• As we saw, experiments show that every C-C bond in benzene is the same length, like $1\tfrac{1}{2}$ bonds, not alternating single and double. This proves the electrons are spread out.

• In terms of orbitals, the six carbon atoms in benzene each provide a 'p-orbital'. These six p-orbitals join up to form a continuous ring of electrons above and below the hexagon. This electron spreading (delocalisation) is what makes benzene so stable and gives it its 'aromatic' properties, often explained by "Hückel's rule" ($(4n+2)$ pi electrons in a special kind of ring).

When we draw benzene showing resonance, we draw a regular hexagon with a circle inside. The circle represents the spread-out pi electrons.

6. Why Aromatic Systems Are Important in Real Life

Many common molecules we know have benzene-like rings and stable electron systems due to resonance. They are everywhere in nature and in things we use daily:

• Vanillin (gives vanilla its flavor)

• TNT (an explosive; resonance in its nitro groups makes it explosive)

• Aspirin (pain reliever; has a benzene ring)

• Menthol (mint flavor; not aromatic itself, but many other flavors are)

• THC (the active part of cannabis)

• Mescaline (a psychedelic substance)

• Some herbicides and their toxic byproducts (dioxins)

• Adrenaline and other brain chemicals (neurotransmitters).

Resonance significantly affects how stable a molecule is (making it less reactive or allowing specific reactions), how it absorbs light (seen in spectroscopy), how drugs work in our bodies (pharmacology), and sometimes how toxic something is. This is all due to how electrons are distributed.

6.1 Neurotransmitter Structures

Many important brain chemicals, like adrenaline, dopamine, and serotonin, have these aromatic rings. Their ability to do their jobs (like controlling mood or our 'fight-or-flight' response) depends on their exact shape and how their electrons are arranged, which is shaped by resonance. The spread-out electrons help them fit perfectly into and interact with 'receptor sites' in our brains.

Understanding resonance is key for designing new drugs, identifying illegal substances, and developing medicines or pesticides. It helps us target molecules precisely.

7. Resonance in Other Common Groups (Not Just Aromatic)

Resonance isn't only for aromatic systems. It's very important for stabilizing many other common groups and short-lived molecules:

7.1 Carbonate Anion $(\mathrm{CO_3^{2-}})$

• The carbonate ion has a central carbon atom connected to three oxygen atoms. Because electrons are spread out, all three C-O bonds end up being equal, even though we draw one as a double bond and two as single bonds in each contributor. Each bond is an average of $1\tfrac{1}{3}$ bonds.

• There are three ways to draw this, simply by rotating where the double bond and negative charges are placed among the three oxygen atoms.

• This spreading out of the $2-$ charge makes the carbonate ion much more stable. This explains why things like calcium carbonate are solid and relatively unreactive.

7.2 Nitro Group $(\mathrm{-NO_2})$

• The nitro group (found in many organic molecules) has a nitrogen atom connected to two oxygen atoms. It has two main resonance contributors: one with an N=O double bond and an N-O$^-$ single bond, and the other with the double bond and negative charge on the other oxygen. So, each N-O bond is effectively $1\tfrac{1}{2}$ bonds.

• In each drawing, the nitrogen has a $+1$ charge, and one oxygen has a $-1$ charge. The whole nitro group is neutral overall.

• When a nitro group is on a benzene ring, the resonance can pull electrons out of the ring onto the nitro group. This makes the ring 'electron-poor'. This pulling of electrons is why nitro groups contribute to the explosive power of compounds like TNT.

7.3 Carboxylate Anion $(\mathrm{RCO_2^-})$

• This forms when a carboxylic acid (like vinegar) loses a hydrogen atom. Carboxylic acids are much 'stronger' acids than alcohols (they lose their hydrogen more easily) because the resulting carboxylate ion is super stable due to resonance.

• The negative charge and the pi electrons are spread out over both oxygen atoms and the carbon. This makes both C-O bonds identical, each like $1\tfrac{1}{2}$ bonds. We can draw two contributors where the negative charge is on one oxygen and then the other.

• This spreading of the negative charge over two electronegative oxygen atoms greatly increases the stability of the ion after the acid loses a hydrogen. This is why carboxylic acids are much more acidic than alcohols, where the negative charge is stuck on just one oxygen. This concept is vital in biology, like in how fats are metabolized or how proteins are structured.

8. How to Calculate Bond Order with Resonance

Simple definition: Bond order is just the number of shared electron pairs between two atoms in a single Lewis structure. For molecules with resonance, where electrons are spread out, we calculate an average bond order to show that bonds are somewhere in between a single and a double bond:

$\text{Bond order} = \frac{\text{Total number of electron pairs shared between two specific atoms (X and Y) across ALL valid resonance contributors}}{\text{Number of equivalent X–Y links in the molecule that show sharing}}$

Example – Carbonate ion ($\mathrm{CO_3^{2-}}$) calculation:

In its three resonance drawings, there's always one C=O double bond (2 shared pairs) and two C-O single bonds (1 shared pair each) between the central carbon and the three oxygens. If we look at one specific C-O bond (say, C-O1), it is shown as a double bond once and as a single bond twice across the three contributors. So, the total shared pairs for C-O1 across all contributors is $2 + 1 + 1 = 4$.

Since there are 3 equivalent C-O bonds in the molecule, the bond order for each C-O bond in the real hybrid is $\frac{4 \text{ (total shared electron pairs for ONE C-O link across all contributors)}}{3 \text{ (number of equivalent C-O links)}} = 1\tfrac{1}{3}$ .

These fractional bond orders calculated from resonance match up really well with the intermediate bond lengths and strengths we find in experiments. This is strong proof that electrons are indeed delocalised.

9. Formal Charge – What It Is & How to Use It

Formal charge is a way to account for electrons in a molecule by pretending that all shared electrons are split equally between the two atoms sharing them. It helps us decide which Lewis structures are the most stable and realistic ones.

Formula (for an atom in a molecule):

$\text{Formal charge (FC)} = \bigl(\text{valence electrons of free atom}\bigr) - \bigl(\text{non-bonding electrons (lone pairs)} + \tfrac{1}{2}\times\text{bonding electrons (shared bonds)}\bigr)$

Simple rules for using formal charge to check resonance drawings:

• The sum of all formal charges on all atoms in a molecule or ion must add up to the total charge of the molecule or ion. This is a quick way to check your math.

• The most stable resonance contributors usually have the fewest non-zero formal charges, or the charges are spread out as little as possible. Lots of separated charges make a molecule less stable.

• If charges can't be avoided, a negative charge is better on an atom that strongly attracts electrons (like oxygen). A positive charge is better on an atom that doesn't attract electrons as much.

Examples from notes:

• Carbon in carbonate drawings has 0 formal charge.

• Each oxygen with only a single bond in carbonate has a $-1$ formal charge. The oxygen with a double bond has 0 formal charge.

• Nitrogen in the nitro group has a $+1$ formal charge, and one oxygen has a $-1$ charge in each drawing.

Understanding formal charge helps chemists guess how molecules will react. Areas with a negative formal charge tend to be 'electron-rich' and like to give electrons, while areas with a positive formal charge tend to be 'electron-poor' and like to accept electrons.

10. Curved Arrows (Showing Electron Movement)

Curved arrows are super important in organic chemistry to show how electron pairs move around during resonance and chemical reactions. They clearly illustrate which bonds are breaking and forming.

• A double-headed arrow ($\leftrightarrow$) between structures means they are resonance contributors (different ways to draw the same molecule, not molecules changing into each other).

• Curved arrows (with a full arrowhead ↷, not half-headed 'fishhooks') on a single structure show electron pair movement:

- The **tail** of the arrow *always* starts from where electrons are: either a lone pair on an atom or a pi bond (the extra bond in a double or triple bond).

- The **head** of the arrow *always* points to where the electron pair is going: either to an atom (to become a new lone pair or help form a new bond) or between two atoms (to form a new pi bond).

- Common patterns: a lone pair forms a new pi bond, a pi bond moves to become a lone pair, or a pi bond moves to form a new pi bond (often in a chain reaction).

• Here's the most important rule: Never move atoms (nuclei) with curved arrows; only electrons move. Breaking this rule means you're confusing resonance with isomers. Also, make sure you don't give second-row elements like carbon, nitrogen, oxygen, and fluorine more than eight electrons in their outer shell (violating the octet rule).

11. How Resonance Affects Reactivity & Stability

Resonance, by spreading out electron density, greatly affects how much energy a molecule has, how stable it is, and how it reacts.

• Resonance makes molecules and short-lived reaction parts more stable (it lowers their energy). This means they are easier to form and directly affects how reactions happen, as reactions tend to follow the easiest (lowest energy) paths. For example, some 'carbocations' (molecules with a positive charge on carbon) are much more stable if their positive charge can be spread out by resonance.

• Resonance explains why some acids are stronger: If an acid loses a hydrogen atom, and the resulting negative charge on the remaining molecule ('conjugate base') can be spread out by resonance, that makes the conjugate base more stable. A more stable conjugate base means the original acid loses its hydrogen more easily, making it a stronger acid. For instance, carboxylic acids are much stronger acids than alcohols because the 'carboxylate anion' (what's left after a carboxylic acid loses H) is resonance-stabilized.

• Electron-withdrawing resonance (pulling electrons away) effects: Groups like the nitro group ($\mathrm{-NO_2}$) or carbonyl group ($\mathrm{-C=O}$) pull electron density away from other parts of the molecule (like aromatic rings) through resonance. This makes aromatic rings less willing to react with 'electrophiles' (electron-loving species) because the ring becomes less electron-rich. They also affect where reactions happen on the ring.

• Electron-donating resonance (pushing electrons in) effects: Groups like those with lone pairs (e.g., an oxygen with lone pairs attached to a ring) can push electron density into an unsaturated system via resonance. This makes aromatic rings more reactive to electrophiles by making them more electron-rich. They also affect where on the ring the reaction will occur.

12. Naming Organic Compounds (IUPAC Rules)

Learning how to name organic compounds correctly (following IUPAC rules) is super important for clear communication in chemistry.

12.1 Basic Structure of a Name

An IUPAC name is like a chemical 'address' that systematically describes the molecule:

Name = Prefix(es) – Parent (root) – Suffix

1. Parent (root) – This part tells you the length of the longest continuous carbon chain (or ring) that includes the most important 'functional group' (a specific group of atoms that gives the molecule its properties). For rings, the ring itself is often the parent.

2. Suffix – This tells you the main functional group in the molecule (the one with the highest priority). If there's more than one functional group, only the most important one becomes the suffix, and the others are named as prefixes.

3. Prefix(es) – These describe all the other parts attached to the main chain/ring, such as smaller carbon groups ('alkyl groups'), halogens (like chlorine), and any 'minor' functional groups. Each prefix has a 'locant' (a number) indicating its position on the chain/ring.

12.2 Step-by-Step Naming Procedure

1. Find and count the longest carbon chain (or ring) that includes the main functional group. If there's a tie in length, pick the chain with the most attached groups (substituents).

2. Identify the highest-priority functional group; this will be your suffix. Number the parent chain from the end that gives this main functional group the lowest possible number. ( บาง functional groups like aldehydes always get C1 by default.)

3. Number all other attached groups (substituents) on the parent chain/ring so that the overall set of numbers for all substituents is as low as possible. Look for the first point where numbers differ.

4. List the names of the attached groups alphabetically, ignoring any numerical prefixes like 'di-', 'tri-', etc. (e.g., 'dimethyl' is alphabetized under 'm'). Each group needs its number(s) before its name.

5. Put the name together: Prefix(es) – Parent – Suffix. Use commas between numbers (e.g., 2,3) and hyphens between numbers and letters (e.g., 2-methyl). There are no spaces in the final name.

Examples: 2-butanol, 3-methylpentane, cyclohexanecarboxylic acid.

12.3 Functional Group Priority Order (Some Examples)

This list tells you which functional group is most important (gets the suffix). Higher numbers mean lower priority, so they become prefixes instead.

1. Carboxylic acid ($\mathrm{-COOH}$) $o$ -oic acid (or -carboxylic acid if it's on a ring).

2. Ester ($\mathrm{-COOR}$) $o$ -oate.

3. Amide ($\mathrm{-CONH_2}$) $o$ -amide.

4. Nitrile ($\mathrm{-C≡N}$) $o$ -nitrile.

5. Aldehyde ($\mathrm{-CHO}$) $o$ -al (or -carbaldehyde for rings).

6. Ketone (\mathrm{>C=O}) $o$ -one.

7. Alcohol ($\mathrm{-OH}$) $o$ -ol.

8. Amine ($\mathrm{-NH<em>2, -NHR, -NR</em>2}$) $o$ -amine.

9. Alkene ($\mathrm{C=C}$) $o$ -ene.

10. Alkyne ($\mathrm{C≡C}$) $o$ -yne.

11. Alkane ($\mathrm{C-C}$) $o$ -ane.

12. Halide ($\mathrm{-F, -Cl, -Br, -I}$), Ether ($\mathrm{-OR}$), Alkyl groups ($\mathrm{-R}$) $o$ always prefixes (halo-, alkoxy-, alkyl-).

12.4 Parent Chain Names for 1-12 Carbons

Prefixes to indicate the number of carbons in the main chain:

12.5 Common Prefixes for Attached Groups

These are common carbon groups when they are attached to the main chain:

• $\mathrm{-CH_3}$: methyl

• $\mathrm{-CH<em>2CH</em>3}$: ethyl

• $\mathrm{-CH<em>2CH</em>2CH_3}$: propyl

• $\mathrm{-CH<em>2CH</em>2CH<em>2CH</em>3}$: butyl

• Other common types: isopropyl, tert-butyl, sec-butyl, isobutyl (important to know the difference).

• Halogens: fluoro-, chloro-, bromo-, iodo-.

• Ethers: methoxy- ($\mathrm{-OCH<em>3}$), ethoxy- ($\mathrm{-OCH</em>2CH_3}$), etc.

• When alcohol is not the main functional group: hydroxy-.

• When amine is not the main functional group: amino-.

12.6 Worked Examples

1. but-3-en-2-one (also known as methyl vinyl ketone)

   • Parent: 'but-' means a 4-carbon chain.

   • Main group: ketone ('-one') is the highest priority, located at carbon #2.

   • Alkene ('-en') is a minor group, starting at carbon #3. Numbering starts from the end that gives the ketone the lowest number (counting right to left).

2. cyclohex-2-enamine

   • Parent: 'cyclohexan-' means a 6-carbon ring.

   • Main group: 'amine' is the highest priority, attached to carbon #1 of the ring (by convention).

   • Alkene ('-en') starts at carbon #2 (meaning the double bond is between C-2 and C-3). The numbering aims to give both the amine and the alkene the lowest possible numbers, with the amine getting C1 if it's the main group.

3. 3-bromo-5-hydroxy-6-aminohept-4-enoic acid (a made-up example with many groups).

   • Parent: 'hept-' (7 carbons).

   • Main group: 'carboxylic acid' ('-oic acid') is the highest priority, always at carbon #1.

   • Minor groups (as prefixes): 'bromo' (at C-3), 'hydroxy' (at C-5), 'amino' (at C-6).

   • Alkene ('-en') is at C-4.

   • Numbering starts from the carboxylic acid carbon (C-1) and goes along the longest chain to give the lowest possible numbers to all the attached groups. When writing the prefixes, put them in alphabetical order (amino, bromo, hydroxy).

13. Practice & Useful Resources

• Chemistry³ textbook – Check out Ch 5.1 (formal charge), Ch 5.5 (resonance), and Ch 2.3 (naming rules).

• Khan Academy videos & exercises are great for visual learners and for checking your understanding of valence electrons, formal charge, drawing resonance structures, and basic naming.

• Angelo State & UIUC (University of Illinois Urbana-Champaign) chemistry handouts online are very good, detailed quick guides for naming complicated organic compounds.

The key to mastering these topics is regular, hands-on practice. Keep drawing all possible resonance structures, calculating formal charges, finding average bond orders, and converting between structures and IUPAC names. This constant practice will truly cement your understanding and skills.

14. Main Takeaways

1. Resonance is a crucial idea that shows how electrons truly behave in many organic molecules – no single drawing tells the whole story. It explains why molecules are extra stable and how they react uniquely.

2. Formal charge is an essential tool for checking if a resonance structure makes sense and how stable it is. It also helps predict where a molecule might react.

3. Bond order in resonance systems is often a fraction (like $1\tfrac{1}{2}$ or $1\tfrac{1}{3}$). This number provides a clear connection to what we see in experiments (like bond lengths).

4. Accurate IUPAC naming requires a step-by-step method: finding the main carbon chain/ring, identifying the most important functional group (which determines the ending of the name), and correctly assigning numbers to all other parts.

5. These basic ideas (resonance, formal charge, and naming) are the foundation for understanding how organic molecules react, how to interpret data from experiments, how drugs work in the body, and the properties of materials. They are critical tools for future chemistry studies and solving real-world problems.

Continual practice and understanding of these principles will make learning higher-level chemistry much easier and provide you with essential tools for solving chemical problems in various scientific fields.