# 3.1 Electronegativity Electronegativity is the measure of the ability of a covalently bonded atom to attract the shared pair of electrons towards itself within a molecule. It is a *dimensionless property* and has no units because it is a relative tendency, not a form of energy. The most commonly used scale for measuring electronegativity was developed by the American chemist **Linus Pauling**. --- ### 3.1.1 Factors Affecting Electronegativity Electronegativity is influenced by several key atomic properties: * **Size of the Atom:** Smaller atoms have a shorter distance between the nucleus and the valence electron shell. This results in a stronger attraction for the shared electrons, leading to higher electronegativity. * **Nuclear Charge:** A higher number of protons in the nucleus results in a greater positive charge. This increased nuclear charge attracts the bonding electrons more strongly, increasing electronegativity. * **Screening by Inner Electrons (Shielding Effect):** Inner-shell electrons repel the valence electrons, "shielding" them from the full attractive force of the nucleus. An increase in the number of inner shells leads to a greater shielding effect, which weakens the nucleus's pull on the shared electrons and decreases electronegativity. In summary: * As **nuclear charge increases** and **atomic size decreases**, the attraction for bonding electrons *increases*, causing electronegativity to *increase*. * As the **shielding effect increases**, the attraction for bonding electrons *decreases*, causing electronegativity to *decrease*. **Table 3.1: Electronegativity of elements in the periodic table** | **Li**: 0.98 | | | | | | | | | | | **B**: 2.04 | **C**: 2.55 | **N**: 3.04 | **O**: 3.44 | **F**: 3.98 | | :----------: | :----------: | :----------: | :----------: | :----------: | :----------: | :----------: | :----------: | :----------: | :----------: | :----------: | :----------: | :----------: | :----------: | :----------: | :----------: | | **Na**: 0.93 | **Mg**: 1.31 | | | | | | | | | | **Al**: 1.61 | **Si**: 1.90 | **P**: 2.19 | **S**: 2.58 | **Cl**: 3.16 | | **K**: 0.82 | **Ca**: 1.00 | **Sc**: 1.36 | **Ti**: 1.54 | **V**: 1.63 | **Cr**: 1.66 | **Mn**: 1.55 | **Fe**: 1.83 | **Co**: 1.88 | **Ni**: 1.91 | **Cu**: 1.90 | **Zn**: 1.65 | **Ga**: 1.81 | **Ge**: 2.01 | **As**: 2.18 | **Se**: 2.55 | | **Rb**: 0.82 | **Sr**: 0.95 | **Y**: 1.22 | **Zr**: 1.33 | **Nb**: 1.60 | **Mo**: 2.16 | **Tc**: 1.90 | **Ru**: 2.20 | **Rh**: 2.28 | **Pd**: 2.20 | **Ag**: 1.93 | **Cd**: 1.69 | **In**: 1.78 | **Sn**: 1.96 | **Sb**: 2.05 | **Te**: 2.10 | | **Cs**: 0.79 | **Ba**: 0.89 | **La**: 1.10 | **Hf**: 1.30 | **Ta**: 1.50 | **W**: 2.36 | **Re**: 1.90 | **Os**: 2.20 | **Ir**: 2.20 | **Pt**: 2.28 | **Au**: 2.54 | **Hg**: 2.00 | **Tl**: 1.62 | | | | | **Fr**: 0.70 | **Ra**: 0.90 | | | | | | | | | | | | | | | ### 3.1.2 Trend of Electronegativity in the Periodic Table #### Down a Group Electronegativity **decreases** from top to bottom within a group. * **Reason:** As you move down a group, a new electron shell is added with each period. This increases the atomic radius and the shielding effect. The increased distance and shielding overcome the effect of the increased nuclear charge, weakening the nucleus's attraction for the shared electrons. *Example: Halogen Group (Group 17)* The electronegativity decreases from Fluorine (F) to Astatine (At). #### Across a Period Electronegativity **increases** from left to right across a period. * **Reason:** As you move across a period, protons are added to the nucleus, increasing the *effective nuclear charge*. The electrons are added to the same principal energy level, so the shielding effect remains relatively constant. The stronger nuclear pull on the electrons decreases the atomic size and increases the atom's ability to attract a shared pair of electrons. *Example: Third Period* The electronegativity increases from Sodium (Na) to Chlorine (Cl). ### 3.1.3 Nature of Bond in Terms of Pauling Electronegativity Values The difference in electronegativity ($\Delta EN$) between two bonded atoms determines the nature of the chemical bond. See also Polar and Nonpolar Covalent Bond and Dipole Moment. | $\Delta EN$ Range | Bond Type | Description | | :--- | :--- | :--- | | $\Delta EN > 1.8$ | **Ionic Bond** | One atom effectively donates an electron to the other, forming ions. | | $0.4 < \Delta EN < 1.8$ | **Polar Covalent Bond** | Electrons are shared unequally, creating partial positive ($\delta^+$) and partial negative ($\delta^-$) charges. | | $\Delta EN \le 0.4$ | **Nonpolar Covalent Bond** | Electrons are shared almost equally between the atoms. | *Note: If $\Delta EN = 0$, the bond is purely covalent.* ### 3.1.4 Covalent Character in a Compound The covalent character in a predominantly ionic compound depends on the **polarizing power** of the cation. A cation with a high charge density (high charge, small size) can distort the electron cloud of the anion, pulling it closer and causing the electrons to be shared. This introduces covalent character. * **Rule:** Higher polarizing power of the cation leads to a greater covalent character. * **Trend:** Polarizing power increases with a higher oxidation state. Therefore, covalent character increases from left to right across a period for metal chlorides, for example. *Example:* Aluminum chloride ($\mathrm{AlCl_3}$) is more covalent than magnesium chloride ($\mathrm{MgCl_2}$). * In $\mathrm{AlCl_3}$, the oxidation state of aluminum is $+3$ ($\mathrm{Al^{3+}}$). * In $\mathrm{MgCl_2}$, the oxidation state of magnesium is $+2$ ($\mathrm{Mg^{2+}}$). The $\mathrm{Al^{3+}}$ ion has a higher charge and is smaller than the $\mathrm{Mg^{2+}}$ ion, giving it a much higher polarizing power. It can strongly polarize the chloride ions, leading to significant electron sharing and a more covalent character in the bond. --- ## Worked Examples **Example 3.1:** Consider the following substances and determine their bond types. (i) Oxygen molecule ($\mathrm{O_2}$) (ii) Potassium Chloride ($\mathrm{KCl}$) (iii) Hydrogen Chloride ($\mathrm{HCl}$) **Solution:** **(i) Oxygen Molecule ($\mathrm{O_2}$)** 1. **Given Values:** Electronegativity of Oxygen ($\mathrm{O}$) = 3.44. 2. **Formula:** $\Delta EN = EN_{atom1} - EN_{atom2}$ 3. **Calculation:** $$ \Delta EN = 3.44 - 3.44 = 0 $$ Since the difference is 0, the electrons are shared equally. The bond is a **nonpolar covalent bond**. **(ii) Potassium Chloride ($\mathrm{KCl}$)** 1. **Given Values:** Electronegativity of Potassium ($\mathrm{K}$) = 0.82; Electronegativity of Chlorine ($\mathrm{Cl}$) = 3.16. 2. **Formula:** $\Delta EN = EN_{Cl} - EN_{K}$ 3. **Calculation:** $$ \Delta EN = 3.16 - 0.82 = 2.34 $$ Since the difference (2.34) is greater than 1.8, the bond is an **ionic bond**. **(iii) Hydrogen Chloride ($\mathrm{HCl}$)** 1. **Given Values:** Electronegativity of Hydrogen ($\mathrm{H}$) = 2.20; Electronegativity of Chlorine ($\mathrm{Cl}$) = 3.16. 2. **Formula:** $\Delta EN = EN_{Cl} - EN_{H}$ 3. **Calculation:** $$ \Delta EN = 3.16 - 2.20 = 0.96 $$ Since the difference (0.96) is between 0.4 and 1.8, the bond is a **polar covalent bond**. Chlorine is more electronegative, so it pulls the shared electrons closer, acquiring a partial negative charge ($\delta^-$), while hydrogen acquires a partial positive charge ($\delta^+$). ---