Chemistry Lesson 1
Atomic and Molecular Structure (Grades 9-12)
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
|HOW MANY ELECTRONS CAN BOND?|
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 the electrons in each shell increase in number:
s subshells have 1 orbital
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.
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 of the Elements. Click here.
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.
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 below. 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. As you can see in the diagram, the 4s subshell is filled before the 3d subshell.
The Periodic Table of the Elements 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.
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 outermost shell (valence electrons).
There are three types of bonding between atoms, ionic bonds, covalent bonds, and metallic bonding. Between ionic bonds and covalent bonds, there is a wide range of bond types, but only 2 labels. Imagine a line with equal sharing of bonding electrons on one end; this is pure covalent bonding, as with molecular hydrogen (H2). With covalent bonds, 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 tightly. This is a polar covalent bond. With the polar covalent bond, the bonding electrons circulate on one atom more than the other. Eventually one atom holds onto the bonding electrons all the time; this is a “pure” ionic bond. Sodium chloride has a “pure” ionic bond. In a “pure” ionic bond, 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 a pure covalent bond. When the electronegativity difference is between 0.1 to 1.7, it is a polar covalent bond. When the difference is greater than 1.7, it is an ionic bond.
Covalent bonds may be single, double, or triple. It all depends on the number of pairs of electrons that are shared. Hydrogen (H2) is a good example of a single covalent bond, since it is made up of two identical atoms that share the 2 electrons equally. Oxygen (O2) is an example of a double covalent bond sharing 4 electrons, and nitrogen (N2) is an example of a triple covalent bond sharing 6 electrons.
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 becomes a positively charged sodium ion (Na+), while a chlorine atom gains one electron and becomes a negatively charged chlorine ion (Cl-). The positive-negative attraction holds many atoms together (NaCl) to form a crystal.
Ionic bonds occur between metals and nonmetals on the Periodic Table of the Elements. Turn to your Periodic Table of the Elements and check out Groups (columns) 1, 2 and 13. 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 in the Periodic Table of the Elements. The elements in Group 1, the Alkali Metals, have 1 valance electron available for bonding. Valance electrons are those electrons in the outermost shell of any element. The elements on the left side of the Periodic Table of the Elements 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.
Metallic bonding is the type of bonding that is exhibited by most metals. Metallic bonding is different from ionic bonds and from covalent bonds. For one thing, it isn’t an isolated bond; it is collective in that a ‘sea’ of electrons travel around all atomic nuclei involved in the metallic bonding, not just around one or two atoms.
Another way of looking at this type of bonding is there is more than one atomic nucleus around which more than one electron orbits. Metallic bonding is like an ionic bond in that it involves positively charged cations and negatively charged electrons, and it is like a covalent bond in that the valence electrons are shared by more than one atom. It is the ‘sea’ of moving electrons that gives metals their properties of high conductivity and high luster.