Physical Chemistry Lesson of the Day – What is the Primary Determinant of the Effective Nuclear Charge for Outer Electrons?

Electrons in the inner shells of an atom shield the electrons in the outer shells pretty well from the nuclear charge.  However, electrons in the same shell don’t shield each other very well.  If an electron spends most of its time below another electron, then the first electron can shield the second electron.  However, this is not the case for electrons in the same shell – they repel each other because they are all negatively charged, and they are at roughly the same average distance from the nucleus.

Thus, the difference between

  1. the charge of the nucleus
  2. and the charge of the core electrons

is the primary contributor to the effective nuclear charge that the outer electrons experience.

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Physical Chemistry Lesson of the Day – Effective Nuclear Charge

Much of chemistry concerns the interactions of the outermost electrons between different chemical species, whether they are atoms or molecules.  The properties of these outermost electrons depends in large part to the charge that the protons in the nucleus exerts on them.  Generally speaking, an atom with more protons exerts a larger positive charge.  However, with the exception of hydrogen, this positive charge is always less than the full nuclear charge.  This is due to the negative charge of the electrons in the inner shells, which partially offsets the positive charge from the nucleus.  Thus, the net charge that the nucleus exerts on the outermost electrons – the effective nuclear charge – is less than the charge that the nucleus would exert if there were no inner electrons between them.

The Gold Foil Experiment and The 250-Million-Ton Pea: The Composition of the Atom

This Atom Is Not To Scale

In a recent post about isotopic abundance, I used a prototypical image of a lithium atom to illustrate the basic structure of an atom.  However, the image was deliberately not drawn to scale to make the protons, neutrons, and electrons visible.  Let’s look at the basic composition of the atom to see why, and we owe this understanding to Ernest Rutherford.  First, let’s give some historical background about what motivated Rutherford to conduct this experiment; we first turn to the Plum Pudding Model by J.J. Thomson.

The Plum Pudding Model

Before 1911, the dominant theory of atomic composition was J.J. Thomson‘s “plum pudding” model.  Thomson hypothesized that an atom consisted of electrons as negatively charged particles (the “plums”) “floating” in a “pudding” of positive charge.

plum pudding model

Plum Pudding Model of the Atom

Source: Wikimedia Commons

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