Chemistry 105 – Problem Set 09 – Electron Configurations Brigham Young University

Problem Set 09 – Electron Configurations

Chem 105

1. (a) 3.86 Explain how the electron configurations of the group 2 elements are linked to their location

in the periodic table. The electron configuration of [core]ns

2

that all the group 2 elements have

places them all in the same column in the s block of the periodic table.

(b) 3.87 How do we know from examining the periodic table that the 4s orbital is filled before the 3d

orbitals? As we start from argon core electrons, we move to potassium and calcium, which are

located in the s block of the periodic table. It is not until Sc, Ti, V, etc., that we begin to fill electrons

in to the 3d subshell.

(c) 3.88 Why do so many transition metals form ions with a 2+ charge? In all of the transition

metals, the s shells (4s, 5s, 6s, 7s) are filled. Because the s shells of the transition metal have a higher

principle quantum number (ie [Kr]4d25s2 for Zr), the electrons in the outer s shells are at a slightly

higher energy than the corresponding d shells. Additionally, the s shells carry 2 electrons like the

group 2 metals in the periodic chart and as a result of these facts, the transition metal s shells are

more likely to be reduced to the 2+ charge.

2. (a) Write out the full electron configuration and the condensed electron configuration for the

following atoms or ions:

Atomic number 15: 1s

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