Origins of the Electronegativity Concept
The idea of electronegativity, or the tendency of an atom to attract bonding electrons, emerged in the early 20th century alongside Lewis‘s groundbreaking vision of the electron pair bond. However, it wasn‘t until 1932 that Linus Pauling formally developed the concept and proposed the first quantitative scale of electronegativity values across the periodic table.
Pauling realized that electronegativity differences between bonded atoms could explain variations in bond properties. By analyzing thermochemical data on covalent bonds, he devised relative numerical values for the electronegativity of each element. This allowed, for the first time, the ability to numerically predict bond polarities.
Over the next decades, several other electronegativity scales were developed, but Pauling‘s scale remains the most widely used due to its simplicity and accuracy. Today, the concept of electronegativity permeates all of chemistry and underlies our modern understanding of chemical bonding.
Defining Electronegativity
Electronegativity refers to the power of an atom to attract shared electrons towards itself when bonded to another atom. The greater an atom‘s electronegativity, the more it draws these electrons.
As introduced by Pauling, electronegativity cannot be directly measured – rather, it is an inherently relative property, with values depending on the atoms being compared. Generally, electronegativity increases moving up and to the right within a period and group, respectively.
This trend mirrors the effective nuclear charge felt by valence electrons. Electronegativity reaches a maximum with fluorine, the most electronegative element, which defines the top end of Pauling‘s scale at 4.0. Elemental francium at 0.7 is the least electronegative element.
Factors That Determine Electronegativity
The key factors that control an element‘s electronegativity are:
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Nuclear charge – This refers to the number of protons in an atom‘s nucleus. More protons exert greater coulombic forces of attraction on electrons. As such, electronegativity tends to increase as we move rightward across a period.
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Distance of valence electrons from nucleus – Electronegativity decreases down a group, because the distance between valence electrons and the nucleus increases due to additional electron shells. This distance weakens the nucleus‘ attraction.
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Electron-electron interactions – Repulsions between valence electrons in the same orbital also play a role. Tighter, more compressed valence electron densities found in lighter elements resist electron attraction by the nucleus.
How to Read an Electronegativity Chart
Electronegativity charts provide a color-coded periodic table denoting Pauling electronegativity values. The trend of increasing electronegativity moving up and rightward is clearly visualized. Values range continuously from 0.7 for francium through to 4.0 for fluorine.
A simple electronegativity chart. Image credit: Jobletter
To predict bonding behaviors, one must:
- Identify the elements present in the bond of interest
- Look up each element‘s electronegativity value
- Subtract the lower value from the higher value
This electronegativity difference (ΔEN) provides insight into the bond type.
Using ΔEN to Predict Bond Type
The electronegativity difference translates to bond polarity and properties:
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ΔEN < 0.5 – Nonpolar covalent bond. Electrons shared equally between atoms. For example, C-H bonds have ΔEN = 0.35, so are nonpolar.
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0.5 < ΔEN < 1.7 – Polar covalent bond. Electrons shared unequally and shift toward more electronegative atom. O-H bonds have ΔEN = 1.38, giving the O δ- charge and H δ+ charge.
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ΔEN > 1.7 – Ionic bond. Electron transferred fully from less electronegative to more electronegative atom, resulting in cation/anion pair. Na+Cl- is an ionic bond with ΔEN = 2.23.
However, one must be cautious using ΔEN alone to determine bond type. Other factors like the metallicity of the elements matters – bonds between metals and nonmetals tend to be ionic. Meanwhile, two nonmetals usually bond covalently.
Let‘s walk through some examples.
Example 1: Predicting Bond Polarity
Consider a C-F bond. Carbon has an electronegativity of 2.55, while fluorine‘s is 4.0.
The difference is: ΔEN = 4.0 – 2.55 = 1.45
Since ΔEN lies between 0.5-1.7, we predict a polar covalent C-F bond, with electron density shifted towards the highly electronegative fluorine atom.
This matches experimental observation. Organofluorine compounds tend to exhibit significant dipoles due to the strongly polar C-F bonds.
Example 2: Explaining Bond Type Variations
Now let‘s analyze HF and HCl:
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HF: ΔEN = 4.0 – 2.2 = 1.8
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HCl: ΔEN = 3.0 – 2.2 = 0.8
Despite HF‘s electronegativity difference exceeding 1.7, it does not form an ionic H+F- bond. Instead, it remains a strongly polar covalent bond. Why is this?
HF consists of two nonmetals – hydrogen and fluorine. Nonmetals preferentially form covalent bonds instead of ionic bonds. Consequently, the ΔEN cutoff values should not be strictly applied in HF‘s case.
Meanwhile for HCl, consisting of the nonmetal hydrogen and halogen chloride, its ΔEN rightly predicts a polar covalent bond.
This demonstrates why simply relying on ΔEN can fail or have exceptions depending on the exact elements present.
Using Electronegativity to Predict Properties
Beyond insights into bond type, electronegativity differences can also forecast bulk properties of compounds:
Polarity – Larger net ΔEN predicts greater molecular dipole moments and polarity. For instance, HF‘s dipole moment is 1.823 D, while HCl‘s is 1.084 D.
Melting and boiling points – Ionic substances with high ΔEN have stronger intermolecular electrostatic forces, increasing boiling/melting points. NaCl (ionic) has a boiling point of 1413°C, versus CO2 (nonpolar covalent) at -78.5°C.
Solubility – Polar molecules with some covalent character are well-solvated by polar solvents like water. Thus, HCl is water-soluble, but nonpolar I2 is not.
Therefore, considering a compound‘s constituent electronegativity differences gives us good initial estimates of its properties.
Limitations and Exceptions
Despite its broad utility, using electronegativity differences to predict chemical bonds and properties has some key limitations:
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It can fail formolecules with multiple identical bonds, like O2, where lone pairs can complicate polarity.
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Some bonds defy the typical ΔEN cutoffs for predicting ionic/polar/nonpolar character. NH3 is a salient example.
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Electronegativity becomes less predictive for larger, more complex organic molecules. Resonance, induction, and orbital hybridization effects take precedence.
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Properties like solubility and intermolecular forces depend on more than just polarity and electronegativity. Molecular geometry and size play a huge role but aren‘t captured by this simple analysis.
In light of these restrictions around electronegativity‘s predictive power, it is best applied for small, binary compounds. Using electronegativity as an initial guide – supplemented by molecular geometry and structure considerations – gives reasonably accurate bond and property predictions.
The Takeaway
At its core, electronegativity remains an immensely useful concept for rationalizing and forecasting the behaviors of different compounds. Simple charts provide reference electronegativity values that can be subtracted to yield electronegativity differences. These differences serve as easy-to-calculate predictors of bond type, polarity, and select physical properties.
Some constraints around electronegativity‘s predictive ability do exist, especially as molecule size and complexity increases. Nevertheless, analyzing electronegativity is an essential first step in the journey towards comprehending chemical bonds on a molecular scale. Combined with a nuanced understanding of geometry and resonance, reliable predictions can be made on how novel compounds may behave.
