Structural Factors in Atoms

In the pages of this section we'll consider several important aspects of atomic structure that influence the periodic atomic properties and the chemical properties of the different elements. Your understanding of these aspects will serve as your foundation for understanding how an element's position on the periodic table is related to its atomic and chemical properties.

Nuclear Charge

The first of these, and the simplest, is the nuclear charge (sometimes referred to as the "actual nuclear charge"). The nuclear charge is the total charge of all the protons in the nucleus. It has the same value as the atomic number. The nuclear charge increases you go across the periodic table. As you get to the end of one period and you go to the beginning of the next period, the atomic number and the nuclear charge continues to increase. Consequently, there is no periodic or repeating nature to the nuclear charge. The nuclear charge just keeps increasing. This is true whether you go across the period or down a group. The symbol for nuclear charge is Z (because it is the atomic number).

Electron Energy Levels

The number of energy levels occupid by electrons in an atom is another very important consideration. As the atomic number increases, so does the number of electrons. But that does not necessarily increase the number of energy levels.

As you go across a period all of the new electrons fit into the energy levels that are already being used. For example, looking at carbon, nitrogen, and oxygen, the number of electrons increases from 6 to 7 to 8. However, when we look at their electron arrangements, notice that all of the electrons are in the first two energy levels.

It is only when you go from one period to the next that you have to increase the number of energy levels. As we go from fluorine to neon to sodium, the number of electrons increases from 9 to 10 to 11. When it passes 10, the third energy level is needed to hold the final electron. So the number of energy levels climbs when we move to a new row.

Next, let’s consider what happens within a group, taking Group IVA as an example. As you go from carbon to silicon to germanium, the number of electrons increases in large jumps, and each jump adds a new energy level, as you can see at right.

Valence Electrons

The valence electrons of an atom are the electrons in the last shell or energy level (note: not just the last subshell. For germanium above, the valence electrons include both the 4s and 4p subshells).

Valence electrons show a repeating or periodic pattern. The valence electrons increase in number as you go across a period. Then when you start the new period, the number drops back down to one and starts increasing again.

For example, when you go across the table from carbon to nitrogen to oxygen, the number of valence electrons increases from 4 to 5 to 6. As we go from fluorine to neon to sodium, the number of valence electrons increases from 7 to 8 and then drops down to 1 when we start the new period with sodium. Within a group - for example, starting with carbon and going down to silicon and germanium - the number of valence electrons stays the same.

A quick way to determine the number of valence electrons for a representative element is to look at which group is it in. Elements in group IA have 1 valence electron. Elements in group IIA have 2 valence electrons. Can you guess how many valence electrons elements in group VIA have? If you guessed 6 valence electrons, then you are correct! The main exception to this rule is helium (He). Despite being in group VIIIA, it has only 2 valence electrons.

Generally speaking, the number of valence electrons stays the same as you go up or down a group, but they increase as you go from left to right across the periodic table. The preceding statement works very well for the representative elements, but it comes a bit short of the truth when you start talking about the transition elements.

We will not consider transition metals in our discussion of valence electrons. The overlap of orbital energies makes it a less useful concept. So we will only consider valence electrons in the context of the representative elements (Groups IA - VIIIA).

Shielding Electrons

Shielding electrons are the electrons in the energy levels between the nucleus and the valence electrons. They are called "shielding" electrons because they "shield" the valence electrons from the force of attraction exerted by the positive charge in the nucleus. They may also be called "core" electrons.

Briefly, what do we mean when we say inner electrons "shield" valence electrons from the nucleus? Electrons are all negative, so they repel one another. This means that, while valence electrons are pulled inward by the nucleus, they are simultaneously pushed outwards by the shielding electrons. This effect isn't enough to push them out of the nucleus entirely, but it does weaken the attraction they feel from the protons in the nucleus.

In fluorine there are 9 protons in the nucleus and there are 2 shielding electrons in the first level between the nucleus and the outer shell. They shield some of the charge of the nucleus from the electrons that are in the outermost energy level.

With sodium, we have 3 energy levels. There is one valence electron in the third level and all the electrons between that one valence electron and the nucleus are shielding electrons. In this case there are 2 in the first energy level and 8 in the second for a total of 10 shielding electrons.

We can see from this that the number of shielding electrons increases when you reach the end of the periodic table and go on to the next period.

Now look at carbon, nitrogen, and oxygen to see that within a period there is no change in the number of shielding electrons. Even though the valence electrons increase in number from 4 to 5 to 6, the number of shielding electrons stays the same - two shielding electrons for each of those elements.

What happens when you deal with the changes within a group? Going from carbon to silicon to germanium, the number of protons in the nucleus increases from 6 to 14 to 32, the number of energy levels increases from 2 to 3 to 4, the number of shielding electrons also increases. In carbon there are 4 valence electrons and 2 shielding electrons. Silicon also has 4 valence electrons, but it has 10 shielding electrons. Germanium (Ge) also has 4 valence electrons, and it has 3 shells or energy levels of electrons that are shielding electrons. There are 2 in the first, 8 in the second, and 18 in the third for a total of 28 shielding electrons along with the 4 valence electrons.

Notice that the shielding electrons follow a pattern somewhat like the number of energy levels. They stay the same within a period. They increase in steps as you start a new period or go down a group .

Effective Nuclear Charge

The next thing to be considered is effective nuclear charge. Generally speaking, effective nuclear charge is the charge felt by the valence electrons after you have taken into account the number of shielding electrons that surround the nucleus.

Again let's take a look at a fluorine atom.

The nucleus itself has a +9 charge and anything in its vicinity will feel that charge. The two electrons in the first energy level as they look at the nucleus feel a +9 charge because that is the charge on the nucleus. But the electrons that are in the valence energy level would be shielded from the nucleus by the 2 shielding electrons. The +9 nuclear charge is shielded by 2 electrons to give an effective nuclear charge of +7 that is felt by the valence electrons.

The symbol for effective nuclear charge is Z_eff, as it is a modification of Z. When dealing with the valence electrons of an atom, we calculate Z_eff by subtracting the number of shielding electrons from the nuclear charge. So, above for fluorine, this looks like:

Z_eff = Z - S

Z_eff = 9 - 2

Z_eff = +7

Let's calculate the effective nuclear charge for the atoms neon, magnesium, and germanium. You should hopefully get values of +8, +2, and +4.

In all these examples the effective nuclear charge is the same as the group number (using the roman numeral system). Get out your periodic table and confirm this is the case. This will be true as long as you are dealing with neutral atoms. Whatever the group number, that will be the effective nuclear charge that the valence electrons feel. It has to be that way for neutral atoms. It is not true when dealing with ions.

Also notice that the effective nuclear charge depends on both the nuclear charge and the number of shielding electrons. The nuclear charge keeps increasing. Meanwhile, the shielding electrons stay constant while you are going across s and p parts of the period. Then when you go to the next period, they jump in number. Consequently, the effective nuclear charge drops at that point. Therefore, the effective nuclear charge increases as you go across a period and then drops and starts over again at +1 when you start the next period. Within a period the effective nuclear charge increases as you go across the periodic table.

As you go down a group, the increase in the nuclear charge is cancelled out by the increase in shielding electrons and the effective nuclear charge stays pretty much the same. In carbon the 4 valence electrons in the outermost shell feel a +6 charge surrounded by two shielding electrons for a +4 effective nuclear charge. For silicon it would also be a +4 effective nuclear charge because the 14 protons in the nucleus are surrounded by 10 shielding electrons. Germanium (Ge) has 32 protons and it has 28 shielding electrons and so the valence electrons feel an effective nuclear charge of +4. As you go down a group, the increase in the nuclear charge is balanced by an increase in the number of shielding electrons so that the effective nuclear charge remains the same.

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