Electronegativity

Electronegativity is a numerical measure of an atom’s tendency to attract electrons toward itself. It captures in a single number what the electron configuration implies qualitatively: elements with nearly full valence subshells hold their outer electrons tightly and hunger for more, while elements with only a few valence electrons hold those electrons loosely and give them up readily. The Pauling scale assigns electronegativity values on an arbitrary but consistent basis: lithium (Li) is assigned 1.0, carbon (C) 2.5, and fluorine (F) 4.0. ^[1] Fluorine, at the top of the scale, is the most electronegative element: it pulls electron density toward itself more strongly than any other. Lithium, near the bottom, gives up its single outer electron with very little resistance.

Atoms possessing low electronegativity easily shed their outermost valence electrons to become cations, whereas those displaying high electronegativity strongly attract additional electrons, driving them to form anions. ^[1] This is why the two properties - electronegativity and ion type - track each other so consistently across the periodic table. Metals sit at the low-electronegativity end and form cations; non-metals sit at the high end and form anions. Electronegativity gives this a quantitative basis: you do not need to know the orbital configuration of every element to predict its behaviour - a single number is sufficient.

The Pauling Scale Values

The electronegativity values for the first 38 elements on the Pauling scale are listed below. These values are used to estimate the degree of ionic character in a chemical bond between any two elements. ^[1]

The noble gases (He, Ne, Ar, Kr) have no electronegativity values because they do not form chemical bonds under normal conditions. Two electronegativity values are given for Cu and Fe because the electronegativity differs between their common oxidation states.

Application to Mineral Bonding

The electronegativity difference between two bonding elements is the key variable in predicting bond character. When two elements with similar electronegativities bond, neither can consistently pull electron density toward itself, so electrons are shared - the bond is covalent. When one element is much more electronegative than the other, it pulls the shared electrons so strongly that effective electron transfer occurs - the bond becomes ionic. Most real mineral bonds fall between these extremes.

Of the eight most abundant elements in the Earth’s crust, oxygen has the highest electronegativity at 3.5. The other seven (Si, Al, Fe, Ca, Na, K, Mg) all form cations when bonding with oxygen and have electronegativities ranging from 1.8 for silicon down to 0.8 for potassium. ^[1] Subtracting the cation values from oxygen’s 3.5 gives electronegativity differences spanning 1.7 for silicon to 2.7 for potassium. ^[1] Si-O bonds, with the smallest difference, have the most covalent character of the common mineral-forming bonds. K-O bonds, with the largest difference, are the most ionic. This gradient is one of the reasons why Si-O tetrahedra are the stiff, directionally bonded building blocks of silicate structures, while K in feldspars and micas is relatively loosely held and readily substituted.

Although several alternative electronegativity scales have been devised, they generally produce comparable relative values. ^[1] The Pauling values are sufficient for the qualitative bond-character assessments that arise most often in mineralogy.

Quantifying Ionic Character

The electronegativity difference between two elements can be converted to an approximate percent ionic character using the empirical relationship: ^[1]

Percent ionic character = (1 - e^-0.25(Δx)2) × 100

For the common oxide-forming cations, this formula translates the electronegativity differences directly into bond character fractions. Bonding in common minerals ranges from about 50% ionic for Si-O bonds (electronegativity difference 1.7) up to about 80% ionic for K-O bonds (difference 2.7). ^[1] This means even the most covalent of the common oxide bonds (Si-O) is half ionic in character - which is why the ionic sphere model of Pauling’s Rules predicts silicate structures as accurately as it does despite silicates being formally classed as covalent in their Si-O tetrahedra.

Ideal ionic bonding is not actually achievable, because the electronegativity difference between any real pair of elements cannot be infinite. The theoretical maximum ionic character is about 93%. ^[1] Even the Na-Cl bond in halite - the textbook example of an ionic bond - is only about 67% ionic, not purely ionic as is sometimes assumed. ^[1] This serves as a reminder that all mineral bonds exist on a continuum: no natural bond is purely ionic, and describing a bond as “ionic” is always a simplification that captures the dominant character without implying an ideal end member.

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