The previous lesson made it clear that the periodic table is based on electron structure.
The elements in the different colored regions have similar properties. The regions indicate elements that have incomplete s,p,d or f subshells. The elements within a colored region are similar, but those in a column within a colored region are even more so. For example, the first two columns contain elements with incomplete s subshells, but the elements of the first column, the alkali metals, are more similar to each other than the elements of the second column, the alkaline earth metals. Similarly the elements of the second to last column, the halogens are more similar to each other than other elements with incomplete p subshells.
To better make sense of this trend we need to revisit the idea of core and valence electrons in light of what we have learned about electron configurations. The idea of core and valence electrons is grounded in the fact that noble gases are very stable. Once formed they tend not to react further to build more complex objects. Noble gases have a unique electron structure. They have filled s and p subshells, the shells in the highest energy level. This leads us to say that atoms tend to change until they achieve the stability of a noble gas electron configuration – one in which the outer shell is complete. These changes can lead to various type of chemical bonds.
Lets make sense of this in the context of a general discussion of how atoms bond. When a proton with a positive charge approaches an electron with a negative charge they bond to form a hydrogen atom. Unlike electrical charges attract to form the atom.
As an equation we would write p+ + e- → H. In general, we think of electrons and protons as building blocks to make atoms.
Once the hydrogen atom forms they no longer have a net charge. Can those neutral atoms make further bonds? That is, can hydrogen atoms bond to form hydrogen molecules? Yes.
As an equation we would write H + H → H2
You might feel puzzled as to how two electrically neutral atoms can attract each. Assuming that the charges cancel each other, there would be no net charge to yield an attractive force to act as a bond. However, if we accept that the charges are not in a fixed position, we can imagine an arrangement yielding more net attraction than repulsion as shown in the simple diagram.
Given that two neutral hydrogens can bond, it is natural to ask if a third neutral atom can join the pair. The answer is no. We do not find any empirical support for the formation of H3.
How about two neutral helium atoms? Will they bond? No.
Here is a summary of these considerations.
How do we account for this? What leads an object with two protons and two electrons (leaving out of our consideration the neutrons), to behave differently than an object with one proton and one electron? In both cases the electrical charge is balanced, so something else must account for the difference.
A clue comes from an attempt to bond two neutral lithium atoms. They form a bond.
Having studied the electron configuration of atoms a natural explanation of this fact emerges. We know that electrons in an atom exist in shells that can only accommodate a fixed number of electrons. When there are two electrons in the first energy level it cannot accommodate any extra electrons. Helium atoms have a complete outer shell. When a second helium atom approaches there is no room in that shell to set up an arrangement leading to a net attraction of subatomic particles. Hydrogen atoms on the other hand do not have a complete outer shell. When a second hydrogen atom approaches there is room in that shell to set up an arrangement leading to a net attraction of subatomic particles. What is true for hydrogen is again true for lithium, where there is room in the second energy level for more electrons.
When based on just a few examples like those above, this reasoning can seem a bit sloppy. For example, in the above we have focused on how electrons might be attracted to incomplete shells, but have not mentioned the loss of electrons from those shells. Nonetheless the ideas will in general prove very useful. When neutral atoms approach each other, their electrons will often shift about until they have achieved a noble gas configuration. These shifts set up distributions of unlike electric charge that allow the atoms to bond to each other. The shifting about has a constraint. Noble gas configurations are stable. We have direct evidence of this from experiment (noble gases tend to be unreactive), and note that all noble gases have a completed outer shell of electrons. Apart from helium which has a configuration of 1s2, all noble gases have completed s and p subshells with a configuration ns2np6, where n stands for the energy level.
In the last chapter we often wrote electron configurations in the condensed form. For example, instead of writing the configuration for sodium as 1s22s22p63s1, we simple wrote [Ne] 3s1. When focusing on the chemical properties of sodium this makes sense. Sodium has 11 electrons, but of these, 10 are in a stable noble gas configuration. The chemical behavior of sodium will involve that single electron that is not part of the noble gas core. We call that single electron a valence electron and the other ten, core electrons. A simple way of showing this is to write sodium as . It is called a Lewis dot or an electron-dot structure. The dot refers to the reactive valence electron and the symbol Na to the nucleus plus all of the stable core electrons. Other atoms can be represented in the same way. For example, chlorine has an electron configuration of 1s22s22p63s23p5. In condensed form it is [Ne]3s23p5, and as a Lewis dot structure,
Here is a portion of the periodic table showing the elements as Lewis dot structures. Since behavior is based on number of valence electrons it starts making sense why the elements in a vertical column behave similarly and thus are given group names.