Electronegativity, Bond Polarity, and Molecular Polarity: Study Notes for Introductory Chemistry

Electronegativity and Bond Polarity

Electronegativity

Electronegativity is a fundamental chemical property that describes an atom’s ability to attract electrons in a chemical bond. It helps predict how electrons are distributed in molecules and the type of bond formed between atoms.

• Definition: Electronegativity is a measure of how strongly an atom attracts electrons in a chemical bond.

• Trend: Electronegativity generally increases across a period (left to right) and decreases down a group (top to bottom) in the periodic table.

• Example: Fluorine (F) has the highest electronegativity (4.0), while cesium (Cs) and francium (Fr) have among the lowest.

Element	Electronegativity
C	2.5
N	3.0
O	3.5
F	4.0
Cl	3.0
Br	2.8
I	2.5

Covalent Bonds

Nonpolar Covalent Bonds

A nonpolar covalent bond occurs between nonmetals with equal or nearly equal sharing of electrons. This happens when the electronegativity difference between the atoms is very small.

• Definition: Electrons are shared equally between atoms.

• Electronegativity Difference: Usually less than 0.4.

• Example: $\text{Cl}_2$ (Cl–Cl): $3.0 - 3.0 = 0.0$ (Nonpolar covalent)

Atoms	Electronegativity Difference	Type of Bond
Cl–Cl	0.0	Nonpolar covalent
Br–Br	0.2	Nonpolar covalent
Si–Si	0.3	Nonpolar covalent

Polar Covalent Bonds

A polar covalent bond forms between atoms with different electronegativities, resulting in unequal sharing of electrons. The electrons are pulled toward the more electronegative atom, creating a dipole.

• Definition: Electrons are unequally shared; the more electronegative atom attracts electrons more strongly.

• Electronegativity Difference: Typically between 0.4 and 1.7.

• Dipole: The bond has a partial negative end ($\delta^-$) and a partial positive end ($\delta^+$).

• Example: $\text{N–O}$: $3.5 - 3.0 = 0.5$ (Polar covalent)

Atoms	Electronegativity Difference	Type of Bond
C–O	0.5	Polar covalent
N–O	0.5	Polar covalent
Cl–F	1.0	Polar covalent

Ionic Bonds

An ionic bond forms when the electronegativity difference between atoms is large (usually greater than 1.7), resulting in the transfer of electrons from one atom to another.

• Definition: Electrons are transferred from the less electronegative atom to the more electronegative atom.

• Example: $\text{NaCl}$: $3.0 - 0.9 = 2.1$ (Ionic bond)

Summary Table: Electronegativity and Bond Type

Electronegativity Difference	Bond Type	Electron Sharing
0 units	Nonpolar covalent	Electrons are equally shared.
0.4–1.7 units	Polar covalent	Electrons are unequally shared; they are attracted to the more electronegative element.
>1.7 units	Ionic	Electrons are transferred from the less electronegative element to the more electronegative element.

Determining Bond Type Using Electronegativity

How to Classify Bonds

To determine the type of bond between two atoms, subtract their electronegativities and compare the difference to standard ranges.

Type of Electron Sharing	Electronegativity Difference	Type of Bond	Reason
Shared equally	2.1 – 2.1 = 0.0	Nonpolar covalent	Between 0.0 and 0.4
Shared about equally	3.0 – 2.8 = 0.2	Nonpolar covalent	Between 0.0 and 0.4
Shared unequally	2.8 – 2.1 = 0.7	Polar covalent	Greater than 0.4
Shared unequally	3.0 – 2.1 = 0.9	Polar covalent	Greater than 0.4
Electron transfer	3.0 – 0.9 = 2.1	Ionic	Greater than 1.7
Electron transfer	3.5 – 1.2 = 2.3	Ionic	Greater than 1.7

Practice: Bond Type Identification

Sample Problems

Use electronegativity differences to classify bonds as nonpolar covalent (NP), polar covalent (P), or ionic (I).

Bond	Difference	Type of Bond
NaCl	2.2	Ionic (I)
H–O	0.5	Polar covalent (P)
Cl–Cl	0.0	Nonpolar covalent (NP)
H–Cl	0.9	Polar covalent (P)

Molecular Polarity

Determining Molecular Polarity

The polarity of a molecule depends on both the polarity of its individual bonds and its overall shape. This determines whether the molecule has a net dipole moment.

• Nonpolar molecules:

  • Contain only nonpolar bonds, or

  • Have polar bonds arranged symmetrically so that bond dipoles cancel.

• Polar molecules:

  • Contain at least one polar bond, and

  • Bond dipoles do not cancel due to the molecule’s shape.

• Example: Carbon dioxide ($\text{CO}_2$) is nonpolar because its linear shape causes the bond dipoles to cancel. Water ($\text{H}_2\text{O}$) is polar because its bent shape prevents dipole cancellation.

*Additional info: The notes above are expanded with definitions, examples, and tables for clarity and completeness, suitable for introductory college chemistry students.*