Lesson 6: Atomic Structure
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. (These are also listed for you in objective 8.) 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 | Main Electron Energy Levels | Valence Electrons
Shielding Electrons | Effective Nuclear Charge
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.
Main Electron Energy Levels
The number of main energy levels for electrons 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 protons increases from 6 to 7 to 8;
so does the number of electrons. However, when we look at their electron arrangements,
notice that all of the electrons are in the first two energy levels. The number of
energy levels being used does not change even though the number of electrons
does. |
|
||||||||||||||||
(This is also shown in example 8a in your workbook. The format there shows the symbol for the element, the number of protons in the nucleus, the electron configuration, and also the total number of electrons in each energy level or shell.) |
|||||||||||||||||
It is only when you go from one period to the next that
you have to increase the number of energy levels. (Also shown in a
different way in part b of example 8 in your workbook.) As we go from fluorine to neon to
sodium, the number of protons increases from 9 to 10 to 11 and thus the number of
electrons increases from 9 to 10 to 11. Notice what happens to the number of energy levels that must be used. For both fluorine and neon, two
energy levels accommodate all of the electrons. But once there are 10 electrons in those
two energy levels (2 in the first and 8 in the second as with neon), any additional
electrons have to go into the next energy level. |
|
Next, let’s consider what happens within a group. (Also shown in example 8c in your workbook.) As you go from carbon to silicon to germanium, the number of protons increases in large jumps. The number of electrons also increase and the number of energy levels used also increases. Notice that carbon has two energy levels, silicon has three, and germanium (Ge) has four levels being used. |
|
||||||||||||||||
The number of energy levels used to accommodate the electrons in the atoms of a particular element is going to be the same as the number of the period. As you go down the periodic table, you will increase the number of energy levels being used. |
|||||||||||||||||
Like the atomic number, the number of energy levels is not really a periodic or repeating feature of atoms. It stays the same throughout a period and then increases by one when a new period starts.
Valence Electrons
The valence electrons are the electrons in the last shell or energy level of an atom. They do 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 only group of representative elements that this method doesn't work for is group 0. Those elements certainly have more than 0 valence electrons; in fact, all of them except for helium have 8 valence electrons. Why doesn't helium have 8 valence electrons? Think for a moment about how many electrons helium has - it has a total of only two electrons, so helium only has 2 valence electrons.
Therefore, 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.
Electrons going into the d sublevels of the transition metals complicate this pattern. In some ways these electrons behave like valence electrons. In some other ways they behave like shielding electrons, which are discussed in the next section. The first electrons into a d sublevel seem to behave more like valence electrons but the last ones seem to act more like shielding electrons, with variations along the way. Switching the order from 4s3d to 3d4s is one way to represent this.
| Sc | Ti | V | Cr | Mn | Fe | Co | Ni | Cu | Zn | |
| outer configuration | 4s23d1 | 4s23d2 | 4s23d3 | 4s13d5 | 4s23d5 |
4s2 3d6 |
4s2 3d7 |
4s2 3d8 |
3d104s1 | 3d104s2 |
| apparent valence electrons | 3 | 2-4 | 2-5 | 2-6 | 2-7 | 2 or 3 | 2 or 3 | 2 or 3 | 1 or 2 | 2 |
As it turns out, the idea of valence electrons is not very useful for transition metals, at least not in a reliable, predictable way.
Electron Dot Diagrams
For a chemist, the valence electrons are quite possibly the most important electrons an atom has. "Why the valence electrons?", you might ask. Well, since the valence electrons are the electrons in the highest energy level, they are the most exposed of all the electrons ... and, consequently, they are the electrons that get most involved in chemical reactions. Chemists use a notation called electron dot diagrams, also known as Lewis diagrams, to show how many valence electrons a particular element has. An electron dot diagram consists of the element's symbol surrounded by dots that represent the valence electrons. The dots are drawn as if there is a square surrounding the element symbol with up to two dots per side. (An element will never have more than eight valence electrons.)
As we discussed above, you can determine how many valence electrons an element has by determining which group it is in. What would the dot diagram for helium look like? It has 2 valence electrons, so it should have 2 dots like this: .He. or He:
Example 6 in your workbook has a few more dot diagrams to study, then try your hand at the ones in Example 7.
Answers for Example 7:
K is in group Ia, so it has 1 valence electron (1 dot).
Al is in group IIIa, so it has 3 valence electrons (dots).
As is in group Va, so it has 5 valence electrons (dots).
F is in group VIIa, so it has 7 valence electrons (dots).
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.
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. (Also look at example 8b in your workbook.) |
|
||||||||||
Next, neon also has 2 shielding electrons along with 8 valence electrons. |
|
||||||||||
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. |
|
||||||||||
Notice 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. (Ex. 8a) |
|
What happen when you deal with the changes within a group? (Also see part c of example 8.) 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 (except for increasing gradually and erratically across the transition metals). 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. (Back to example 8b in your workbook.) 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. If you get out beyond the valence electrons, then the effective charge is 0 simply because the +9 charge of the nucleus is surrounded by 9 electrons. |
|
|||||||||||||||||
Generally we are only concerned with the effective nuclear charge as it pertains to the valence electrons (+7 in this case), but sometimes the broader concept of what charge is felt by other electrons is useful.
Practice
With that in mind, figure out what the effective nuclear charge would be for neon and sodium for the electrons in each energy level. What would the first two electrons feel? What would the next eight electrons feel? Then for sodium, what would that last electron feel? So take a moment to figure that out. (Refer to ex. 8b. in your workbook, if that helps.)
|
|
|
|
|
Answers
|
|
|
|
|
As you go across the table from fluorine to neon, the effective nuclear charge felt by the valence electrons increases. Then as you go to sodium in the next period, there is another energy level, the number of shielding electrons increases, causing the effective nuclear charge felt by the valence electron to drop. |
|
|||||||||
In all these examples the effective nuclear charge is the same as the number of valence electrons. That is true as long as you are dealing with neutral atoms. However many valence electrons there are, 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, (but increase gradually across the d part 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. |
|
One other thing I should mention about the effective nuclear charge is that it is quite often referred to as the kernel charge. The "kernel" includes the nucleus and all shielding electrons but does not include the valence electrons.