Periodic trends: Electronegativity

Periodic trends: Electronegativity

Lessons

In this lesson, we will learn:
  • To understand the definition of electronegativity and how it is measured.
  • To apply understanding of electrostatic principles to the trends in electronegativity.
  • To predict the electronegativity of elements in comparison to each other.

Notes:

  • As seen in Periodic trends: Atomic radius, chemists have found, through experimenting, some principles of electrostatic forces – forces that exist because charged particles attract or repel each other. Some principles are:
    • #1: Oppositely charged particles attract each other, while particles of like charge repel each other.
    • #2: The greater the charge difference of two particles, the greater their force of attraction (for example, the attractive force between a 2+ ion and a 2- ion is stronger than the attractive force between a 1+ and a 1- ion).
    • #3: Attractive forces between oppositely charge particles decrease with distance.
    • #4: Repulsive forces between like-charged particles decrease with distance.

  • Electronegativity is the ability of an atom (specifically the nucleus) to attract bonding electrons to its outer electron shell. It is measured using the Pauling scale – fluorine is highest at 4.0 on the scale, the most electronegative element, whilst francium is the lowest at 0.7 and is the least electronegative element.

  • Electrostatic theory can explain the trend in electronegativity in the Periodic Table, both across a period and down a group:
    • As you go across the elements in a period, each element has an extra proton in its nucleus, increasing its nuclear charge.
      • This increased nuclear charge means the nucleus attracts its outer shell electrons with increasingly greater force and also attracts other electrons able to complete the outer shell if it isn’t full.
      • Therefore, as you go across the period (where nuclear charge increases and electron shielding doesn’t change), it is easier for atoms to attract bonding electrons into their outer shell. This means electronegativity is higher.
    • As you go down a group of the periodic table, each element has an extra shell of electrons as well as increased nuclear charge.
      • The extra inner shell of electrons shields (by repulsion) the outer shell electrons from the positively charged nucleus.
      • This means with more electron shells, the nucleus does not attract electrons to fill its outer shell as easily. As a result, electronegativity decreases as you go down a group in the Periodic Table.
    • Remember: Noble gases are not given an electronegativity value because their atoms generally do not form bonds and since they already have full outer shells, they do not attract electrons to complete full outer shells!

  • The trends in electronegativity mean fluorine is the most electronegative element. The effect of electron shielding down a group is more influential than the effect of increased nuclear charge across a period, so oxygen is the second most electronegative element (around 3.5 on the Pauling scale), followed by chlorine (around 3.0).

  • The difference in electronegativity of atoms affects how different atoms bond with one another and can lead to substances of varying properties. This is a very important part of chemistry which the next chapter will look at – bonding between atoms and the properties of compounds they make as a result!>.
  • Introduction
    Summary of electronegativity
    a)
    Definition of electronegativity.

    b)
    Trends in electronegativity in the Periodic Table – using electrostatic principles to explain them.

    c)
    What happens when two atoms have different electronegativity?


  • 1.
    Recall and explain the trends in electronegativity in the periodic table.
    Describe and explain the trend in electronegativity across the elements in period 2 of the Periodic Table.

  • 2.
    Recall and explain the trends in electronegativity in the periodic table.
    Describe and explain the trend in electronegativity in group 2 of the Periodic Table.