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Chapter 5: Problem 1

What is electronegativity? What are the periodic trends in electronegativity?

Short Answer

Expert verified
Electronegativity is the ability of an atom to attract bonding electrons, and it increases across a period and decreases down a group in the periodic table. Noble gases are exceptions as they do not typically form bonds that would be used to assign electronegativity.

Step by step solution

01

Understanding Electronegativity

Electronegativity is a measure of how strongly an atom can attract a bonding pair of electrons towards itself when it is part of a compound. The most common scale for electronegativity is the Pauling scale, named after the chemist Linus Pauling who introduced it.
02

Periodic Trend: Across a Period

In the periodic table, electronegativity increases from left to right within a period. This is because atoms have more protons (positive charge) as one moves across a period, so the ability to attract electrons improves due to the increased nuclear charge.
03

Periodic Trend: Down a Group

Electronegativity decreases as one goes down a group in the periodic table. This is because the atomic radius increases, which means that the bonding electrons are further from the nucleus and are less strongly attracted.
04

Exceptions to the Trend

The noble gases are usually not assigned electronegativity values because they typically do not form compounds since they already have complete valence electron shells. Thus, they do not compete for electrons in bonding.

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Key Concepts

These are the key concepts you need to understand to accurately answer the question.

Pauling Scale
In chemistry, understanding how atoms interact is crucial, and one key aspect is their electron affinity or electronegativity. The Pauling scale offers a way to quantify this trait. Devised by Linus Pauling, the scale helps predict how strongly atoms will attract electrons in a chemical bond.

Measured on a relative scale where the value for hydrogen is set at 2.20, other elements are compared against this standard. For example, fluorine, the most electronegative element, has a Pauling value of 3.98. This scale is indispensable for predicting the behavior of atoms in molecules, determining whether bonds will be ionic, covalent, or somewhere in between.
Periodic Trends in Electronegativity
Electronegativity varies in a predictable way across the periodic table. As you move from left to right across a period (row), the electronegativity typically increases. This happens because atoms gain more protons in their nuclei, enhancing their ability to draw electrons.

However, as one moves down a group (column), electronegativity decreases. The added electron shells increase the atomic radius, distancing the valence electrons from the nucleus, thus weakening the pull. This periodic trend is vital for understanding how different elements will bond and interact.
Atomic Radius
The atomic radius is the distance from an atom's nucleus to the outermost shell of electrons. It has an inverse relationship with electronegativity. As the radius increases, generally when moving down a group, it's harder for the nucleus to attract electrons to form a chemical bond:

  • Larger atomic radius = weaker pull on bonding electrons
  • Smaller atomic radius = stronger pull on bonding electrons
Knowing an element's atomic radius can help predict its reactivity and the types of bonds it will form.
Nuclear Charge
The nuclear charge is the total charge of the protons in an atom's nucleus and directly affects electronegativity. Higher nuclear charge can lead to a stronger attraction for bonding electrons if the atomic radius doesn't increase significantly. This concept explains why electronegativity increases from left to right across a period:

  • More protons in the nucleus = higher nuclear charge
  • Higher nuclear charge (without significantly larger radius) = stronger attraction for electrons
This is why elements towards the right of a period (excluding noble gases that don't form many compounds) are more electronegative.

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Most popular questions from this chapter

Draw the Lewis structure for urea, H2NCONH2, one of the compounds responsible for the smell of urine. (The central carbon atom is bonded to both nitrogen atoms and to the oxygen atom.) Does urea contain polar bonds? Which bond in urea is most polar?

Write the Lewis structure for each molecule or ion. a. H3COH b. OH- c. BrO- d. O2 2-

In the very first chapter of this book, we described the scientific approach and put a special emphasis on scientific models or theories. In this chapter, we looked carefully at the Lewis model of chemical bonding. Why is this theory successful? What are some of the limitations of the theory?

Draw the Lewis symbols for the atoms Al and O. Use the Lewis model to determine the formula for the compound that forms from these two atoms.

List the correct electron and molecular geometries that correspond to each set of electron groups around the central atom of a molecule. a. four electron groups overall; three bonding groups and one lone pair b. four electron groups overall; two bonding groups and two lone pairs c. five electron groups overall; four bonding groups and one lone pair d. five electron groups overall; three bonding groups and two lone pairs e. five electron groups overall; two bonding groups and three lone pairs f. six electron groups overall; five bonding groups and one lone pair g. six electron groups overall; four bonding groups and two lone pairs
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