3 Radiation & Atomic Theory+
3.6 Quantum Numbers
3.9 Ionic Bonding
3.11 Covalent Bonding
When two non-metal elements get together we have a dilemma in a way. BOTH elements by there nature "want" more electrons to achieve noble gas electron configuration - they want 8 electrons in their outer shell as we say which is s2p6. So neither element is going to give up electrons, they are holding on to what they got (thank you Bon Jovi). Or maybe even they're saying to their electrons "Never Gonna Give You Up" as they Rickroll into a covalent bond. The POINT being... they have no problem in sharing electrons so that they can each make it to 8. An example is carbon tetrachloride (CCl4 - we'll get to naming in a bit). The carbon has 4 electrons in its outer shell (2s22p2) and each of the four chlorines has 7 (2s22p5). So what happens is that each chlorine shares one electron with the carbon and the carbon shared one electron with each of the chlorines (that's 4 total). The end result is four covalent bonds between the carbon and the chlorines. All five atoms are holding on tightly to the shared electrons - this is the basis of all covalent bonds. This is shown below with electron dot formulas. Each dot is a valence electron around that atom.
Notice how the finished product has 8 electrons (dots) around each element. The diagram below emphasizes this fact by circling the 8 electrons around each atom. The other diagram/drawing shows how we swap out a shared pair of dots for a line. The line represents 2 shared electrons and we call it a covalent bond.
This is the way we diagram a molecule with covalent bonds. When you start learning this skill, you will definitely use all the dots and move things around until you find the structure with all 8's surrounding the atoms. We will dive into electron line/dot formulas a little later, for now lets learn some simple naming for binary covalent compounds.
So a covalent bond is based on the sharing of one or more electron pairs between two non-metal atoms. A "perfect" covalent bond means the sharing is perfectly equal - meaning each atom participating in the bond has an equal share of the electrons. This is basically a 50/50 split of electrons. However, this can only happen perfectly when the two atoms have identical electronegativities (EN values). So to be truly perfectly covalent an atom needs to bond with itself, then you are guaranteed that each atom pulls on the electrons the same amount.
So anytime the two atoms are different, there is most likely a difference in EN values and therefore an unequal sharing of the electrons. This is the basis of polar covalent bonds. A polar covalent bond is when one of the atoms gets a bit more of the electrons - technically meaning an unequal sharing of the electron pair. The more electronegative atom will pull the electrons to itself a bit more than the other atom. This leads to a slight partial negative charge (δ–) on the more electronegative atom and a partial positive charge (δ+) on the more electropositive atom. This is a polar bond. Polar bonds can add up on a molecule to give a polar molecule which has a net dipole. Please also note that polar bonds can also just cancel each other out to result in a non-polar molecule as well. You really have to know the three dimensional shape of a molecule plus all the polarities of the bonds to determine if a molecule ends up polar or non-polar. Here is a line/dot structure for HCl (hydrogen chloride) that also shows the partial charges on the H and the Cl.
Chlorine is more electronegative than hydrogen which means it will pull the electrons more toward itself and away from hydrogen. The result is a partial negative charge on the chlorine and a partial positive charge on the hydrogen. Anytime there is a net separation of positive and negative charge for a molecule, the molecule is polar and will have a net dipole moment which is just a measure of the partial charge separation.
What I mean by binary covalent compounds is that only two different elements make the compound. It is a bit like ionic compounds except instead of a cation and an anion, you have element1 and element2 and those elements are non-metals.
Who's first? In a similar way to ionic compounds (cations are first), you should always list the more electropositive element first - that is the one with the lower electronegativity value (another reason to learn and memorize the trend).
Prefixes: For covalent compounds we will have to use prefixes to tell others how many of an atom there is in the compound. A very simple example of this is carbon monoxide (CO) and carbon dioxide (CO2). The mono- prefix means one and the di- prefix means two. Can you gues what SF3 is? Sulfur trifluoride. Notice we name the second element (which is always the more electronegative one) as an -ide, just like we did for monatomic anions. Learn your prefixes so you can get the counts right.
mono-, di-, tri-, tetra-, penta-, hexa-, hepta-, octa-, nona-, and deca- will cover you from 1 to 10 which is plenty.
Avoid Double Vowels (sometimes): When using a prefix ending in o or a for the number (which is all of them except di- and tri-) you might need to drop the o or a when you combine it with oxide. Examples are to write and say carbon monoxide, and do NOT say or write carbon monooxide (note the double vowel). N2O5 is dinitrogen pentoxide (not pentaoxide). This avoid double vowels is mainly for fixing the names of oxides. You DO keep a double vowel for something with iodide. So CI4* is carbon tetraiodide (not tetriodide). So fix it on oxides and avoid the ao and oo double vowels.
* Look out for the san-serif font thing. That is a capital C and a capital i, I. Here's a lowercase ell, l and an uppercase I. Our exams are printed in a serif font so it is obvious: CI4.
A friendly reminder: Do NOT get confused and start using your prefix knowledge with ionic compounds. Remember, ionic compounds just name the ions. Covalent compounds name the elements where the first is the element and the second is the -ide version AND we use prefixes for counts.