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Rule and Principle of Electrons Help (page 2)

By — McGraw-Hill Professional
Updated on Aug 28, 2011

Examples

Example 1

Using the period and group information of the Periodic Table, what is the configuration of phosphorus (P), atomic number Z = 15?

Start with the first subshell on the Periodic Table as 1s, then in the second period (row) you have 2s. Jumping across in the same row is 2p. In the third period, there is 3s, 3p. In the fourth period there is 4s, 3d, and 4p.

1s 2 (first period) 2s 2 2p 6 (second period) 3s 2 3p 3 (third period) Phosphorus is in period 3 so n = 3 and group 5A so valence electrons = 5 (Remember, the valence electron shell arrangement is the same as the outermost placement of electrons.)

Example 2

Try nickel (Ni), atomic number Z = 28. First consult the Periodic Table for the period number and group. Then start building the subshells. Did you get 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 8 ? This can also be written with the subshells grouped together as 1s 2 2s 2 p 6 3s 2 p 6 d 8 4s 2 .

Example 3

If cesium (Ce) has a configuration of 1s 2 2s 2 p 6 3s 2 p 6 d 10 4s 2 p 6 d 10 5s 2 p 6 6s 1 what is its period number and group? What is the filling valence subshell?

Cesium is an alkali metal in period 6 and group 1. The filling orbital is 6s 1 .

It is important to remember that electrons of different elements in the same groups look the same with the exception of different orbital configurations. Members of the same family have identical valence electron structures, when considering numbers of paired and unpaired valence electrons. These similar valence structures make it possible for family group members to react similarly.

Additionally, the elements with the most orbitals, further and further out from the nucleus, become increasingly more reactive as the electrons zip through a larger area. They have more "party room" to come in contact with other atoms and they take advantage of it!

Ionization Energy

The period of an element also describes patterns. If you look at the elements from left to right in the Periodic Table along any row

The ionization energy of an element is the energy needed to detach an electron from an atom of an element.

Regular property changes can be compared to changes in electron arrangement. The higher the number of electrons in the outermost shell of an atom, the higher the ionization energy of that atom.

The elements are listed in rows so that it is easy to find information quickly. For example, if you were interested in a specific element, you would check out its place in the Periodic Table. Who are its neighbors on the chart? Which group is it in? Which period? How many electrons are in its outermost orbit (sometimes called outermost shell)? Is it reactive or not? Is it a metal or non-metal? Look again at Figure 6.4. All the group and period information that you will need on the elements can be found on the Periodic Table.

Knowing the reactivity of an element is important. If the element to be studied was potassium and you put it into water, you would have a wild reaction since alkali metals get really crazy in water. They give off hydrogen gas that ignites with the heat of the reaction and gives off a violet flame, like a lot of mini-fireworks. From potassium's soft solid state and its low boiling point, you might think it is a mild-mannered element. However, the reactive heat from an encounter with water changes a solid chunk of potassium into a liquid by melting it. Can't judge an element by its cover!

Practice problems for these concepts can be found at – Electron Configuration Practice Test

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