Difficulty: Introductory | Prerequisites: Basic atomic structure, periodic table familiarity
Electronegativity is one of the first properties you need to internalise in organic chemistry. It determines how electrons are shared in a bond, which in turn controls whether a bond or molecule is polar. Polarity drives solubility, reactivity, boiling points, and nearly every intermolecular interaction you will study from here on. If you skipped general chemistry or are rusty on periodic trends, start here.
Electronegativity measures how strongly an atom pulls a shared pair of electrons toward itself in a bond. When two bonded atoms have different electronegativities, the bond is polar. Whether the whole molecule is polar depends on its geometry: if the polar bonds do not cancel out symmetrically, the molecule has a net dipole.
Electronegativity
A measure of the tendency of an atom to attract a bonding pair of electrons toward itself. Values are given on the Pauling scale, where fluorine is the highest at 4.0.
In simple terms, this means: the higher the number, the more that atom hogs the electrons in a bond.
Polar bond (polar covalent bond)
A covalent bond in which the electrons are shared unequally because the two atoms have different electronegativities. The more electronegative atom carries a partial negative charge (δ−), the other a partial positive charge (δ+).
Think of it as: a tug-of-war over the electrons where one side is stronger, so the rope sits off-centre.
Nonpolar covalent bond (pure covalent bond)
A covalent bond between atoms with identical or nearly identical electronegativities. The electron pair is shared equally, so no partial charges develop.
In simple terms, this means: a perfectly even tug-of-war.
Ionic bond
A bond formed when the electronegativity difference between two atoms is so large that one atom effectively takes the electron entirely, producing a cation and an anion.
Think of it as: one side yanks the rope away altogether.
Polar molecule
A molecule that has a net dipole moment because its polar bonds do not cancel each other out. Determining polarity requires looking at the molecule's three-dimensional geometry, not just individual bonds.
In simple terms, this means: the molecule is lopsided in how its charge is distributed.
Dipole moment
A quantitative measure of the separation of positive and negative charge in a molecule. A molecule with a dipole moment is polar.
Think of it as: a number that tells you how lopsided the charge distribution is.
The type of bond that forms between two atoms depends on the difference in their electronegativity values.
Pure nonpolar covalent bond: No electronegativity difference (identical atoms). Example: H–H, Cl–Cl.
Polar covalent bond: Small electronegativity difference. Electrons are shared but pulled toward the more electronegative atom. Example: H–Cl, C–O.
Ionic bond: Large electronegativity difference. One atom effectively takes the electron from the other, forming ions. Example: Na–Cl, K–F.
There is no sharp cutoff between these categories. They sit on a spectrum. A difference of roughly 0.4 to 1.7 is typically considered polar covalent, above 1.7 is ionic, and below 0.4 is essentially nonpolar. Your course may use slightly different thresholds.
This distinction trips up many students. They are not the same thing.
Polar bond: Only looks at the electronegativity difference between the two atoms in that specific bond.
Polar molecule: Looks at the molecule's overall geometry and whether the individual bond dipoles cancel or add up.
A molecule can contain polar bonds and still be nonpolar if the geometry is symmetrical enough that the dipoles cancel. CO₂ is the classic example: each C=O bond is polar, but the molecule is linear and symmetric, so the two dipoles point in opposite directions and cancel to zero.
For a molecule to be polar, two conditions must both be met:
The molecule must be "lopsided" in some way, meaning its geometry is asymmetric.
The molecule must have a distinct region of partial positive charge and a distinct region of partial negative charge.
If either condition fails, the molecule is nonpolar even if it contains polar bonds.
Electronegativity follows two clear trends:
Across a period (left to right): Electronegativity increases. Atoms gain more protons, pulling electrons in more tightly.
Down a group (top to bottom): Electronegativity decreases. Atoms get larger, and the bonding electrons sit further from the nucleus with more shielding.
Fluorine (top right of the periodic table, excluding noble gases) has the highest electronegativity of any element. Francium (bottom left) has the lowest.
Noble gases are generally exempt from electronegativity rankings because they are chemically inert and do not typically form bonds.
Picture the periodic table. Draw an arrow pointing right across any period and another arrow pointing upward within any group. Both arrows point toward increasing electronegativity. The convergence point, top right (excluding Group 18, the noble gases), is fluorine.
Polar bonds are drawn with a dipole arrow (→) pointing from the less electronegative atom (δ+) toward the more electronegative atom (δ−). Some textbooks use the crossed-arrow notation, where the tail has a small plus sign.
Electronegativity values are empirical (looked up from a table, not calculated from a formula at this level). The key skill is reading the Pauling scale values and computing the difference between two atoms to classify the bond.
Water is polar because of its bent geometry, and that polarity is the reason it dissolves salts, supports life, and has an unusually high boiling point for its molecular weight. Understanding bond and molecular polarity is the foundation for predicting which solvents dissolve which solutes ("like dissolves like"), why oil and water do not mix, and how drug molecules interact with biological receptors.
Students often think that having polar bonds automatically makes a molecule polar. It does not. Molecular geometry determines whether the bond dipoles cancel. CO₂ has polar bonds but is a nonpolar molecule.
Students sometimes confuse electronegativity with electron affinity. Electronegativity describes an atom's pull on shared electrons within a bond. Electron affinity is the energy change when a free atom gains an electron. They are related but measure different things.
Some students assume ionic and covalent bonds are completely separate categories. They are endpoints on a continuous spectrum. Most real bonds sit somewhere in between.
Students occasionally forget that noble gases are excluded from electronegativity trends because they typically do not form bonds. Seeing them on the periodic table and trying to assign them a place in the trend is a common error.
⚠️ Expect questions asking you to rank atoms or bonds by polarity using electronegativity values.
⚠️ A common exam question gives you a molecular structure and asks whether the molecule is polar or nonpolar. You must consider geometry, not just individual bonds.
⚠️ Periodic table trends (electronegativity increasing across a period, decreasing down a group) are tested frequently and are often combined with questions on atomic radius, ionisation energy, and electron affinity.
⚠️ Be prepared to classify a bond as nonpolar covalent, polar covalent, or ionic given two elements and their electronegativity values.
True or false: Electronegativity increases as you move down a group on the periodic table. (False. It decreases.)
Fill in the blank: The element with the highest electronegativity is _____. (Fluorine.)
True or false: A molecule with polar bonds is always a polar molecule. (False. Geometry determines whether bond dipoles cancel.)
Fill in the blank: A bond with no electronegativity difference is called a _____ bond. (Pure nonpolar covalent.)
True or false: Noble gases are typically excluded from electronegativity discussions because they are inert. (True.)
Q: Rank the following bonds from least polar to most polar: C–H, O–H, N–H.
A: C–H (smallest electronegativity difference), then N–H, then O–H (largest electronegativity difference). Oxygen is more electronegative than nitrogen, which is more electronegative than carbon.
Q: Carbon dioxide (CO₂) contains two polar C=O bonds. Explain why the molecule is nonpolar.
A: CO₂ is linear. The two C=O bond dipoles point in exactly opposite directions and cancel each other out, giving the molecule a net dipole moment of zero.
Q: Water (H₂O) also contains two polar bonds (O–H). Why is water polar while CO₂ is not?
A: Water has a bent geometry (approximately 104.5° bond angle). The two O–H bond dipoles do not point in opposite directions, so they do not cancel. The molecule has a net dipole moment.
Q: Using periodic trends, predict which atom is more electronegative: oxygen or sulphur. Explain your reasoning.
A: Oxygen. Both are in Group 6A, but oxygen is above sulphur. Electronegativity increases as you move up a group because the bonding electrons are closer to the nucleus with less shielding.
Q: Classify the bond in NaCl as nonpolar covalent, polar covalent, or ionic. Justify your answer.
A: Ionic. The electronegativity difference between sodium (0.93) and chlorine (3.16) is approximately 2.23, which is well above the typical threshold of 1.7 for ionic character.
This connects to molecular geometry and VSEPR theory (Chapter 1.7 and beyond) because you need the three-dimensional shape of a molecule to determine whether it is polar. Without geometry, you can only assess bond polarity, not molecular polarity.
This also connects to intermolecular forces (dipole-dipole interactions, hydrogen bonding) covered later in the course. The strength and type of intermolecular forces depend directly on whether a molecule is polar, which traces back to the electronegativity concepts here.
Acid-base chemistry in organic reactions relies heavily on electronegativity. The ability of a substituent to stabilise a negative charge (and therefore strengthen an acid) is rooted in how electronegative its atoms are.
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