Instruction 1-6 Connection Among the Location in the Table, the Atomic Number, and Mass | How to Identify Metals, Semimetals, Nonmetals, and Halogens | How to Identify Alkaline Metals, Alkaline Earth Metals, and Transition Metals | Lanthanide, Actinide, Transactinide, and Transuranium Elements | Ionization Energy, Electronegativity, Relative Sizes | How Many Electrons Can Bond? | Size and Mass | Location and Quantum Electron Configuration | Summary |
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As we said previously, it is very helpful to
know the exact electron configuration of each element. You were told that
there was
additional restriction on the electrons in the shells. The shells are
divided into subshells and orbitals. The subshells are called s, p, d, f,
g, h, and i. Each subshell has an odd number of orbitals and increases in
number; s subshells have 1 orbital, p subshells have 3 orbitals, d
subshells have 5 orbitals, etc. The figure shows only one orbital from
each subshell. The subshells have different configurations; s subshells
are spherical (like a ball), p subshells have 2 lobes (like dumbbells),
and d subshells have 4 lobes (like a clover leaf). There is a maximum of 2
electrons in each orbital. There are good reviews on these web pages. http://www.orbitals.com/orb/index.html http://www.colby.edu/chemistry/OChem/DEMOS/Orbitals.html We’ll list the electron configuration of the 18 most common elements
here. But for the rest, you’ll need to look at a good Periodic Table.
Click: The notation used below tells us which shell and subshell the electrons are occupying in the lowest energy state. The Pauli Exclusion Principle states that one electron goes into each orbital of a subshell of the same energy, before any orbital gets a second electron.
Here are the maximum number of electrons that can be contained in any one shell, and then divided into subshells: Remember that the atomic number of any element indicates the number of protons in its nucleus. A neutral atom always has the same number of electrons as protons. The number of electrons available for sharing is determined by the number of electrons in the outer most shell (valence electrons). There are three types of bonding between atoms, ionic bonding, polar covalent bonding and pure covalent bonding. Actually, there is a wide range of bonding, but only 3 labels. Imaging a line with equal sharing of bonding electrons on one end; this is pure covalent bonding, as in hydrogen. In pure covalent bonding, the bonding electrons are shared equally between the atoms. As you move to the other end, one of the atoms gets greedy and holds onto the bonding electrons more often. This is polar covalent bonding. In polar covalent bonding, the bonding electrons circulate on one atom more than the other. Eventually one atom holds onto the bonding electrons all the time; this is “pure” ionic bonding. Sodium chloride is almost “pure” ionic bonding. In “pure” ionic bonding, there is no sharing of the bonding electrons.
It is possible to define the transition between the different types of bonding if you look at the electronegativity (from the last instruction) of the atoms. When the difference in electronegativity is zero, it is pure covalent bonding. When the difference is between 0.1 to 1.7, it is polar covalent bonding. When the difference is greater than 1.7, it is ionic bonding. Covalent bonds Atoms can also become stable by losing or gaining electrons to form ions. This kind of bond is called an electrovalent or ionic bond. In an ionic compound, positively charged ions and negatively charged ions stick together; much like magnets sticking together. Ordinary table salt, sodium chloride, is a good example.
To form sodium chloride, a sodium atom (Na) loses one electron and Ionic bonds http://en.wikipedia.org/wiki/Ionic_bond occur between metals and nonmetals on The Periodic Table. Turn to your Table and check out Groups (columns) 1 (IA), 2 (IIA) and 13 (also known as IIIA). These Groups provide many of the positive partners involved in ionic bonding. The easiest way to determine the number of electrons that can bond is to look at the element’s location on the Periodic Table. The elements in Group 1, the Alkali Metals, have 1 valance electron available for bonding. http://en.wikipedia.org/wiki/Valence_electron Valance electrons are those electrons in the outer most shell of any element. The elements on the left side of the table lose electrons to become more stable. So in this case, the Alkali Metals become cations (positive ions) with a +1 charge. The elements in Group 2, the Alkaline Earth Metals, have 2 valance electrons available for bonding; they become cations with a +2 charge. We pass the transition elements because they have a variable number of electrons for bonding. Group 13 have 3 valance electrons available for bonding and become cations with a +3 charge. Group 14 have 4 valance electrons available for bonding; they also have 4 spaces to hold electrons. As we get to the right side of the table, the elements want to hold onto the electrons they have and grab more if it is possible. Group 14 elements are usually covalently bonded, but can either be cations with a +4 charge, or anions with a 4 charge. Group 15 have 5 valance electrons and 3 spaces to hold electrons, which are used for bonding; they become anions with a -3 charge. Group 16 have 6 valance electrons and 2 spaces to hold electrons, which are used for bonding; they become anions with a -2 charge. Group 17, the halogens, have 7 valance electrons and 1 space to hold an electron, which is used for bonding; they become anions with a -1 charge. Group 18, the noble gases, have 8 valance electrons and no space to hold any electrons. This is why the noble gases don’t react.
Now let's do Practice Exercise 1-6 (top).
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When
the electrons are arranged in the atoms, we see an order in which orbitals
are filled. You
can easily create the order as shown in the diagram at right. The order of
filling is 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s,
6f, 5d, 7p. Filling all these subshells will give us 118 elements (2 more
than what has been discovered). As you can see in the diagram, the 4s
subshell is filled before the 3d subshell. The Periodic Table reflects
this subshell filling as well. How you may ask? Look at the figure and see
where each of the subshell is listed. As you should recall, the
Lanthanides and Actinides fit into the table between 6s and 5d and between
7s and 6d. 
