Source: Lecture notes, The Ohio State University
Tags: ionic bonds, covalent bonds, polar covalent, electronegativity, ionisation potential, electron affinity, electrostatic potential maps, dipole, bond polarity, valence electrons
Difficulty: Introductory Prerequisites: Familiarity with atomic structure, electron shells, and the periodic table (general chemistry level).
Before you can understand how organic molecules react, you need to understand how atoms hold together. This section covers the two main bonding types (ionic and covalent), introduces the spectrum of bond polarity in between, and explains the driving forces behind bond formation: ionisation potential, electron affinity, and electronegativity. Most bonds in organic chemistry are covalent, and most of those are polar covalent, so getting comfortable with partial charges and dipoles here pays off throughout the course.
Atoms bond either by transferring electrons (ionic) or sharing them (covalent). Most organic bonds are covalent, and because atoms rarely share electrons perfectly equally, most covalent bonds are polar. Electronegativity differences between bonded atoms determine where on the ionic-to-covalent spectrum a bond sits.
Covalent bond
A bond formed by the sharing of electrons between two atoms. In simple terms, both atoms contribute electrons to a shared pair that holds them together.
Ionic bond
A bond based on the electrostatic attraction of two ions with opposite charges. One atom gives up an electron entirely, and the other takes it.
Ionisation potential (IP)
The energy it takes to remove an electron from an atom. Think of it as how tightly an atom holds onto its electrons. Hydrogen has an IP of 314 kcal/mol (1314 kJ/mol).
Electron affinity (EA)
The energy released when an electron attaches itself to an atom. A large electron affinity means the atom strongly "wants" an extra electron. Hydrogen's EA is approximately 18 kcal/mol (75 kJ/mol).
Electrostatic potential map
A visual representation of how charge distributes itself across a molecule's electron cloud. Blue regions indicate positive charge (low electron density); red regions indicate negative charge (high electron density).
Core configuration
The representation of an atom using letters for the nucleus and dots for the valence (outermost) electrons. Also called electron-dot notation.
Polar covalent bond
A bond where a pair of electrons is unequally shared between two atoms. Most bonds in organic chemistry are polar covalent.
Electronegativity
A measure of an atom's ability to attract shared electrons in a bond. Fluorine is the most electronegative element (4.0 on the Pauling scale).
Electropositive
Describes an atom that tends to donate electrons in a bond (low electronegativity).
Electronegative
Describes an atom that tends to accept or attract electrons in a bond (high electronegativity).
Electric dipole
A separation of opposite charges within a molecule, symbolised by an arrow crossed at its tail (pointing from the positive end toward the negative end).
Delta notation (δ⁺ / δ⁻)
Symbols used to indicate partial positive and partial negative charges on atoms in a polar covalent bond. δ⁻ means partial negative charge; δ⁺ means partial positive charge.
Many atoms bond to carbon in an intermediate way, falling somewhere between purely ionic and purely covalent.
Ionic bonds: one atom fully transfers its electron(s) to another. The result is a cation (positive) and an anion (negative) held together by electrostatic attraction.
Covalent bonds: atoms share electrons rather than transferring them.
Ionisation potential tells you how much energy is needed to pull an electron away from an atom. High IP means the atom holds its electrons tightly.
Electron affinity tells you how much energy is released when an atom gains an electron. High EA means the atom readily accepts electrons.
Hydrogen can either lose an electron to become a bare proton (H⁺, the bare nucleus) or gain one to become a hydride ion (H⁻, helium configuration).
H → H⁺: IP = 314 kcal/mol
H → H⁻: EA ≈ 18 kcal/mol
Formation of ionic bonds between identical elements is difficult because electron transfer is unfavourable when both atoms have similar IP and EA.
The high IP of hydrogen also prevents hydrogen halide bonds (like H–F) from being ionic.
Valence electrons are the outermost electrons, shown as dots around the element symbol.
Period 2 elements and their valence electron counts:
Li (1), Be (2), B (3), C (4), N (5), O (6), F (7)
Na, Mg, Al, Si, P, S, Cl follow the same pattern in Period 3.
H· + ·H → H:H (shared pair between two hydrogens)
H· + ·Cl: → H:Cl: (shared pair, with lone pairs remaining on chlorine)
:Cl· + ·Cl: → :Cl:Cl: (shared pair, with lone pairs on each chlorine)
Pure covalent bonds occur between atoms of equal electronegativity (e.g. Cl–Cl). Electronegativity difference ≈ 0.
Polar covalent bonds occur between atoms of differing electronegativity (e.g. H–Cl). Electronegativity difference between roughly 0.5 and 2.0.
Pure ionic bonds occur between elements at the two extremes of the periodic table (e.g. Na–Cl). Electronegativity difference > 2.0.
The most ionic bonds occur between elements at opposite ends of the table (least covalent character).
The most covalent bonds occur between atoms of equal electronegativity.
H: 2.1
C: 2.5, N: 3.0, O: 3.5, F: 4.0
Li: 1.0, Na: 0.9, K: 0.8 (very electropositive)
Cl: 3.0, Br: 2.8, I: 2.5
When two bonded atoms differ in electronegativity, the bond is polarised: the more electronegative atom carries a partial negative charge (δ⁻) and the less electronegative atom carries a partial positive charge (δ⁺).
This charge separation is called an electric dipole.
In carbon dioxide (O=C=O), the two polar C=O bonds point in opposite directions, so the dipole moments cancel and the molecule has no net polarisation.
Blue = positive charge (electron-poor region).
Red = negative charge (electron-rich region).
These maps show a form of the molecule's electron cloud and help visualise how charge is distributed.
Electronegativity difference and bond type:
0 to ~0.5 → pure covalent
~0.5 to ~2.0 → polar covalent
above ~2.0 → ionic
Polar bonds are why water dissolves salt but not oil, and why soaps work (they have both polar and nonpolar regions). Electronegativity differences drive the polarity of drug molecules, influencing whether they can cross cell membranes or dissolve in blood. Electrostatic potential maps are used in computational chemistry and drug design to predict how molecules will interact with biological targets.
Students often think a bond is either purely ionic or purely covalent. In reality, most bonds sit somewhere on a spectrum between the two, and most organic bonds are polar covalent.
A common mistake is confusing ionisation potential with electron affinity. IP is the energy to remove an electron; EA is the energy released when gaining one. They measure opposite processes.
Students sometimes assume CO₂ is a polar molecule because it has polar C=O bonds. The molecule is actually nonpolar overall, because the two dipoles cancel each other out due to its linear geometry.
Electronegativity values are not the same as electron affinity. Electronegativity describes behaviour within a bond; electron affinity describes an isolated atom gaining an electron.
⚠️ Be able to classify a bond as ionic, polar covalent, or nonpolar covalent based on electronegativity difference.
⚠️ Know the electronegativity trend across a period (increases left to right) and down a group (decreases going down).
⚠️ Understand how to read an electrostatic potential map: blue = positive, red = negative.
⚠️ Be prepared to explain why CO₂ has polar bonds but is a nonpolar molecule (opposing dipoles cancel).
⚠️ Know the difference between IP and EA, and be able to use them to explain why certain bonds are ionic or covalent.
True or false: A bond with an electronegativity difference of 1.5 is best described as polar covalent.
Fill in the blank: On an electrostatic potential map, blue indicates ______ charge.
True or false: The most electronegative element is oxygen.
Fill in the blank: δ⁺ indicates a partial ______ charge.
True or false: CO₂ is a polar molecule.
Answers: 1. True. 2. Positive. 3. False (it is fluorine, at 4.0). 4. Positive. 5. False (the opposing dipoles cancel, making it nonpolar overall).
Q: What is the difference between ionisation potential and electron affinity?
A: Ionisation potential is the energy required to remove an electron from an atom. Electron affinity is the energy released when an atom gains an electron. They describe opposite processes.
Q: Why is the bond in NaCl classified as ionic rather than covalent?
A: Sodium (electronegativity 0.9) and chlorine (3.0) have a large electronegativity difference (2.1), which means the electron is essentially transferred rather than shared.
Q: Explain why CO₂ is nonpolar despite having polar bonds.
A: CO₂ is a linear molecule. The two C=O dipoles point in exactly opposite directions and cancel each other out, resulting in no net dipole moment for the molecule.
Q: What does a red region on an electrostatic potential map indicate?
A: A region of high electron density, carrying a partial or full negative charge.
Q: Where on the periodic table would you find the most electronegative elements?
A: In the upper right corner (excluding noble gases). Fluorine, oxygen, nitrogen, and chlorine are among the most electronegative.
This section builds directly on the periodic table and atomic structure from general chemistry. It connects forward to Lewis structures (Section 1.4), where you will use valence electrons and electronegativity to draw bonding diagrams and assign formal charges. Understanding bond polarity here is also essential for later topics on molecular geometry, intermolecular forces, and reaction mechanisms throughout organic chemistry.
Ionic bonds, covalent bonds, polar covalent bonds, electronegativity, Pauling scale, ionisation potential, electron affinity, electrostatic potential map, dipole, partial charge, delta plus, delta minus, valence electrons, electron-dot notation, bond polarity, nonpolar, NaCl, CO₂, organic chemistry Chapter 1, OSU organic chemistry