MHS Chemistry
Bonding Summary

There are just under one-hundred elements that occur in nature, but there are millions and millions of different compounds. Your hair, the air you breathe, the ink in a pen, the components of paper, the metals and plastics making your computer, and the water in the fountains are just a very few examples of the chemicals that exist everywhere in everything you touch. For us to be able to talk about and control them, we must be able to organize them. And we organize them by what kinds of bonds hold them together. The bonds holding atoms together give any chemical it's very basic general characteristics.

A chemical bond occurs when two atoms are held together by mutual attraction to the same electrons. This attraction is balanced by the repulsion of the nuclei for each other, and the repulsion of the electrons for each other. Since each element is a unique combination of protons and electron arrangements, atoms of each element have a slightly different attraction for electrons. This relative attraction for shared electrons has been tabulated as electronegativity, which is a 0 to 4 scale. The atom with the strongest attraction for shared electrons is fluorine, with an electronegativity of 3.98, and the lowest electronegativity is that of Francium (0.7).  The symbol for electronegativity is c (that's "c" in the symbol font).

The type of bond formed between two atoms depends on their electronegativity. Atoms with strong attractions for electrons are non-metals, and tend to form anions (negative ions). Atoms with weak attractions for electrons are metals, and tend to form cations (positive ions). Atoms with moderate attraction for shared electrons are known as transition metals.   There is one family of atoms that have virtually no attraction for electrons, but also do not give up the ones they have, called the noble gases. [Click here for a review of metals and non-metals and their properties.]

There are some atoms that do not have electronegativities listed. Most of these the atoms are so rare that it has not been possible to gather experimental data. A few of them (the noble gases, see below) do not form any compounds, and so a relative attraction for shared electrons cannot be calculated.

Chemical bonds occur when electrons end up paired with each other, and the bonded atoms always have lower total energy than the separated atoms.

Ionic Bonding
If two atoms have very different attractions for electrons, then one of them will "steal" the electrons from the other. These two atoms are then "stuck" together by their opposite charges, in what is known as an ionic bond. Atoms in ionic compounds do not need exactly opposite charges; for example, calcium chloride has the formula CaCl2 and consists of calcium ions with a 2+ charge and chloride ions with a 1- charge. There will be enough of each ion so the overall charge is zero.

Also, ionic substances always have their ions in specific ratios (like in calcium chloride above: 1 Ca : 2 Cl), but they do not exist as molecules. Instead, they exist as a crystal lattice, which is a regularly constructed arrangement of positive and negative ions. These lattices can be any size, from sub-microscopic to many feet across, but for a given compound, they all have the same chemical properties

Because an ionic compound does not exist as a molecules of a specific size, we cannot calculate a molecular weight.  We may still need to know how much of a certain compound we have, so for ionic compounds we calculate formula mass.  This is calculated and used the same way as molecular weight, but it tells us the mass of a single "formula unit" of a substance.  For table salt, even though there are no specific Na-Cl pairs, we still add up the mass of one sodium atom and one chlorine atom, because they make up the formula unit.

Ionic compounds that dissolve in water and break into their individual ions are known as electrolytes because the resulting solution conducts electricity. Ionic compounds that do not dissolve in water are called non-electrolytes, and tend to involve larger ions. Almost all ionic compounds are solids at room temperature. [If you can think of one that isn't, e-mail me.]

Covalent Bonding
Two atoms with similar strong attactions for electrons can't "steal" them from each other, so they must "share" electrons. This is known as a covalent bond. Generally, atoms will form a covalent bond if they are both non-metals. Also, many metals (except from the first two columns) will form covalent bonds with non-metals, because even though they have opposite tendencies to form ions, they are not so different that complete transfer of an electron will take place.

Although some covalent compounds dissolve in water, like sugar or vinegar, their solutions do not conduct electricity because they do not break into ions. It may seem that this describes a non-electrolyte ionic compound, but there is another difference: covalent compounds usually form molecules, which are the smallest unit of a compound with all the properties of that compound. Unlike lattices, molecules have definite composition. A molecule of water has the formula H2O, which means that there is a tiny unit with exactly two hydrogen atoms connected to exactly one oxygen atom. A different number of either atom would be a different compound with different properties.

Covalent compounds are often liquid or gas at room temperature, and the ones that are solid are often soft or waxy.

Metallic Bonding
We have seen that a given pair of atoms can either both strongly attract electrons (covalent bond), or one can strongly attract electrons away from the other (ionic bond). There is a third possibility that occurs if neither atom involved has a strong attraction for other electrons. These atoms are metals, and the resulting situation is known as a metallic bond. In this case, many atoms will be sharing valence electrons, but so weakly that the electrons do not "belong" to any specific nucleus. The collection of atoms acts like a clump of chocolate chip cookie dough, with each chip being a nucleus, and the dough being the electrons. They are all held together, and they hold a consistent shape, unless you push on them. In that case, the whole system deforms, but the same nuclei and electrons are still there. That's a model for malleability.

We won't be worrying too much about metallic bonds, except to say that two or more metals together form metallic bonds.

How Can We Tell?
There are two ways to decide which type of bond is involved in a given compound. There is a "rule of thumb" method, and a "calculation" method. They're both pretty easy. "Rule of Thumb" means an easy pattern to remember.

The "rule of thumb" method requires you to mentally divide the periodic table into four regions, shown below.
active metal + active metal makes a metallic bond
active metal + transition metal makes a metallic bond
active metal + non-metal makes an ionic bond
transition metal + non-metal makes a covalent or ionic bond
transition metal + transition metal makes a metallic bond
non-metal + non-metal makes a covalent bond
anything + noble gas makes no bond



In other words, the closer together they are on the right of the table, the more likely they are to form covalent bonds.

The calculation method requires a periodic table with electronegativities listed (like the colorful ones you're all supposed to have!). To find out what type of bond two atoms will form, subtract their electrongativities (big minus small). If the difference is bigger than 1.67, it's ionic. If the difference is less, it's covalent or metallic (two metals).


Here's another way of thinking of it.
If two atoms have strong similar attractions for electrons, they will form covalent bonds (two non-metals)
  weak similar attractions for electrons, they will form metallic bonds (two metals)
  very different attractions for electrons, they will form ionic bonds (metal + non-metal)
If any atom has NO attractions for electrons, they will form NO bonds (noble gases)



Try these:

  type of bond using "rule of thumb" method type of bond using calculation method
Li & Cl ________________________ ________________________
C & Cl ________________________ ________________________
Fe & S ________________________ ________________________
Ba & I ________________________ ________________________
N & O ________________________ ________________________
Ga & Br ________________________ ________________________
Fe & Ne ________________________ ________________________
Ru & N ________________________ ________________________
Ni & Cu ________________________ ________________________
Be & S ________________________ ________________________

Did the two methods give different results?

Polarity Of Bonds
When calculating the electronegativity difference to determine whether a given bond is covalent or ionic, it may have occured to you that sometimes the difference may be barely enough to be ionic or covalent. What if the difference is exactly 1.67? Well, it helps to remember that even in the most ionic bond, there is still a little bit of time the "lost" electron spends around the positive ion. As the difference in electronegativity approaches zero, the sharing of the electron becomes more and more even, and the bond becomes less and less ionic.

To clear up these possibilities a little more, we call any bond between two non-metals with the same electronegativity a pure covalent bond. If the atoms have similar (but different) electronegativities, they are said to form a polar covalent bond. In these cases, the electron spends a bit more time closer to the more electronegative atom. For example, in water, the H-O bond is polar, with the oxygen "hogging" the electron. When the hogging is extreme, we have an ionic bond.

So what if the difference is 1.68? Or 1.66? Well, call it what the calculation tells you (ionic for the first one, covalent for the second), or use the rule of thumb, and show how you reached your decision. It is unlikely anyone will argue with your logic.

[][MHS Chem page]