Calculate formal charges on atoms in a compound. Lewis structures, also called Lewis dot diagrams, model covalent bonding between atoms. Specify drawing parameters in the Lewis Structure Drawing Tool. Open the Lewis Structure Drawing Tool, lewisdrawauthoring2013.html, in a browser that has the Flash plugin. Set the grading options for the problem on the lower left: Click a grading setting option to toggle whether or not it will be active.
Ammonia is NH 3. Nitrogen has five valence electrons and each hydrogen brings one to the molecule. It's easy to see that those three electrons from the hydrogens could complete an octet on the nitrogen: 5 + 3 = 8.There's a twist in this molecule—a small one—and it gives ammonia some startling properties, some of which are beyond the scope of these notes, but trust me, they're cool.The figure below shows how to construct the Lewis structure. Start with each atom surrounded by its valence electrons, 5 for N, 1 for H. I like to color each atom's electrons differently so I can keep track of them, but it's not absolutely necessary. Water is a very important molecule, so it's important to understand its bonding, which will in turn create all of its other properties.The Lewis structure of water shows that the oxygen atom has two lone pairs.Those lone pairs, together with the large difference in electronegativity between oxygen and hydrogen, give water one of its most important properties, its strong polarity.In the rendering on the right below you can see that water, which actually is a bent planar molecule (again, you wouldn't necessarily know this just from the Lewis structure), has a negative end and a positive end. More correctly, it has one end that is more negative than the other (called δ-) and one more positive (called δ+).
The unbonded electron pairs create a region of dense negative charge. And because oxygen holds the bonding electrons of the H-atoms tightly to itself, the H-atoms are essentially bare protons hanging off the oxygen.The polarity of water and its ability to hydrogen bond gives water some of the properties that are deeply intertwined with the chemistry of living things on Earth. It's for that reason that it's difficult for us to conceive of life on another planet without water – but you never know. Some Lewis structures will lead to bonding that is ambiguous. A double bond might be present between an atom and one or more other equivalent partners.
Which one to choose? This is called a degeneracy, and it turns out that nature tends to pick both, neither, and a combination of the two bonds.As an example, let's look at the Lewis structure of nitric acid, HNO 3. First the atoms with their valence electrons:Now we can arrange the bonds in two ways. In both, all atoms have a full valence shell. Here they are:Take a minute and convince yourself that every atom (except the hydrogen) has an octet of electrons in its valence. Here are the two structures in stick form:So which one does nature pick? Well, it turns out that whenever we have two equivalent structures like this (we have two degenerate structures or a degeneracy), nature picks a combination of both, and we're better off writing the two bonds more like 1-1/2 bonds, like this:Notice that the red oxygen is different than the other two.
It's bound to a hydrogen and the electronegativity difference makes this bond more ionic than covalent. The result is that the hydrogen can detach as a bare proton quite readily, leaving a NO 3 - ion behind.
That's why HNO 3 is an acid. Now let's see how a molecular ion, the carbonate ion, CO 3 2-, can bond stably. When we work with ions, we begin with the usual number of valence electrons of a neutral molecule, in this case four for carbon and six for each of the three oxygens. But this is a 2- ion, so we'll add two electrons to the neutral mix to give it that -2 net charge.Those last two electrons can be filled in anywhere they're needed to form a full valence.
We start with the raw materials:Now put them together and use the two extra (red) electrons to fill in any gaps in order to form full valence shells for every atom. The sulfate ion, SO 4 2-, is a very interesting exception to many of our assumptions about bonding. In fact, it caused a lot of argument in the chemistry community early on. At first glance, with sulfur and oxygen both holding six valence electrons, we might draw a Lewis structure like this:The problem is that the two extra charges needed to stabilize this molecule are localized on the sulfur atom. Nature tends to spread that extra charge out, and sometimes at a cost that can contradict what we've already learned. In this case, it's the octet rule. In fact, SO 4 2- tends to bond more like this:Now the extra charges are a little more spread out, and we can see that there's nothing special about our double bond locations, which means resonance structures and even more spreading of charge.
But the twelve electrons around the sulfur are troubling. Remember that sulfur has d-electrons, and some of those are used as valence electrons. Overall this structure is more stable than the all singly bonded one.Here are all of the resonance forms of SO 4 2-, and they lead to an ion with four identical bonds that are somewhere between double and single in strength.It was a measure of the S-O bond length in SO 4 2- that led to a deeper investigation of the bonding.You shouldn't feel like you ought to have recognized this case.
It took some Nobel-prize winning chemists, including Linus Pauling, some time and significant argument to uncover the truth. It says a lot about the nature of the electron, if you think about it. The sulfate ion (above) is one case of an exception to the octet rules. There are others.
I'm not sure there's any point in memorizing such exceptions; better to know that they exist and be wary of them. Know some of the signs that an exception might exist. Here are a few examples. Phosphorus pentachloride, PCl 5Phosphorus pentachloride ( PCl 5) is an exception to the octet rule. You can see the Lewis structure of PCl 3 in the practice problems below. Because a chlorine atom only needs one electron to complete its valence shell, it shares one and only one electron with phosphorus, so in PCl 5, phosphorus is surrounded by a total of ten electrons. It does this by using its d-shell electrons.A more 3-dimensional structure is shown on the right.
Three chlorines are in a plane (blue triangle) and the line containing the other two cuts through the center of that triangle and is perpendicular to it. The arrangement is called a trigonal bipyramid. Sulfur hexafluoride, SF 6For similar reasons, sulfur can bind to six fluorine atoms with single bonds in a square bipyramid arrangement.Xenon hexafluoride, XeF 6Finally, we don't normally think of noble gases interacting with anything, let alone forming bonds. But it turns out that if the noble gas is big enough, like xenon (Xe), the d-orbitals can allow bonds to form. Xenon forms a couple of covalent compounds, one of which is XF 6.XF 6 looks just like SF 6, but the central Xe is now surrounded by 14 electrons.
Sometimes it's difficult to tell which of two possible Lewis structures of a compound represents the actual bonding of the molecule. In those cases we resort to calculating what's called the formal charge of each atom. Formal charge is just a way of bookkeeping that helps us to decide which of multiple Lewis structures is the likely true bonding arrangement of a covalent molecule. The sum of the formal charges, with a couple of extra rules, will help us to decide which of multiple-possible valid Lewis structures is likely to be the correct one.
Here's how it's done. Calculating formal chargeFor each atom. Count the number of valence electrons of the neutral atom. Subtract the number of non-bonding electrons (usually in lone pairs). Subtract the number of bonds shared by the atom.Example: CH 4 (methane)The carbon in CH 4 has four electrons as a neutral atom.
It has no lone pairs, and it shares four bonds, so the formal charge is zero. Each hydrogen atom has one electron as a neutral atom, no lone pairs and shares one bond, for a formal charge of zero. All atoms in the molecule have zero formal charge, the 'happiest' situation for any molecule. Example: H 3C(CO)CH 3, (acetone)The central carbon has a formal charge of 4 (valence electrons) - 0 (lone pairs) - 4 (bonds) = 0. The oxygen has a formal charge of 6 - 4 - 2 = 0 (same ordering of terms). Each of the methyl (CH 3) carbons has a formal charge of 4 - 0 - 4 = 0. Distinguishing between two valid Lewis structuresHere is an example of a case where we can find two valid Lewis structures for a compound, fulminic acid (HCNO).
We can use formal charges to decide which is most likely to be the actual arrangement of atoms. Here are the structures:Now let's calculate the formal charges of the lower structure, using double-bonds:Note that the carbon has a formal charge of -1 and the nitrogen a charge of -1.
The formal charges of the structure with the triple bond look like this:Here, the oxygen — the most electronegative element in the molecule — has the negative charge, and the nitrogen retains its +1 charge. This structure is more likely to be the correct one, because the negative charge is on the most electronegative element of C, N and O. Here are the formal charges on each atom for each bonding arrangement. The fully double-bonded structure (right) has the lowest formal charges on each atom.
Even though sulfur has a bonded valence of 12 electrons, this is still the most stable structure. A few elements in the third row of the periodic table, plus a great many elements with d-electrons, are capable of this.This bonding arrangement of SO 3 is confirmed by experiment, which shows that its structure is trigonal-planar (a flat molecule with oxygen atoms at the vertices of an equilateral triangle. VSEPR theory predicts that the 1- and 2-double bonded structures would be a little different (but distinguishable).While formal charge assignment isn't too useful in other areas of chemistry, it can be really enlightening when questions about bonding like this need to be resolved.
Covalent Lewis Dot StructuresA bond is the sharing of 2 electrons.Covalent bonds share electrons in order to form a stable octet around each atom in the molecules. Hydrogen is the exception it only requires 2 electrons (a duet) to be stable.How do we draw a covalent Lewis Dot Structure?Level 1 (basic)1. Add up all the valance electrons of the atoms involved. Ex CF 4So C has 4 and F has 7 (x4 we have 4Fs) = 32 valence electrons2. You need to pick the central atom. This is usually easy, this atom will be surrounded by the others. Never H.So C will be surrounded by F's.3.
Now we create our skeleton structure by placing bonds in. A bond is a dash that represents 2 electrons.We have now placed 8 electrons as 4 bonds. We have 32-8= 24 more to place.4. Starting with the outer atoms add the remaining electrons in pairs until all the electrons have run out.All 32 electrons are now in place, count the dots around each F. 6 dots and a bond (2 electrons) is 8. We have our octet.The carbon has 4 bonds (2electrons) for its 8.DONELevel 2 (Double and Triple bonds)Same rules apply until #41.
Add up all the valance electrons of the atoms involved. Ex CO 2So C has 4 and O has 6 (x2 ) = 16 valence electrons2. You need to pick the central atom. This is usually easy, this atom will be surrounded by the others.
Never H.So C will be surrounded by O's.3. Now we create our skeleton structure by placing bonds in. A bond is a dash that represents 2 electrons.We have now placed 4 electrons as 2 bonds. We have 16-4=12 more to place.4. Starting with the outer atoms add the remaining electrons in pairs until all the electrons have run out.All 16 electrons are now in place, count the dots around each O.
6 dots and a bond (2 electrons) is 8. We have our octet.The carbon has 2 bonds (2electrons) for its 4.?We need 8, so move a pair of electrons from the O to between the C and O. It will share 2 pairs of electrons instead of 1. It now has a double bond instead of a single bond.carbon has 6 electrons, so move 2 from the other oxygennow they all have an octet, it cleans up like thisMake it symmetrical.Level 3-Lewis Dots of Polyatomic IonsSame rules apply, at the end they get brackets and a chargeAP Chemistry and or College Level Rules1. Determine whether the compound is covalent or ionic. If covalent, treat the entire molecule. If ionic, treat each ion separately.
Compounds of low electronegativity metals with high electronegativity nonmetals ( DE N 1.7) are ionic as are compounds of metals with polyatomic anions. For a monoatomic ion, the electronic configuration of the ion represents the correct Lewis structure. For compounds containing complex ions, you must learn to recognize the formulas of cations and anions.2. Determine the total number of valence electrons available to the molecule or ion by:(a) summing the valence electrons of all the atoms in the unit and(b) adding one electron for each net negative charge or subtracting one electron for each net positive charge. Then divide the total number of available electrons by 2 to obtain the number of electron pairs (E.P.) available.3.
Organize the atoms so there is a central atom (usually the least electronegative) surrounded by ligand (outer) atoms. Hydrogen is never the central atom.4. Determine a provisional electron distribution by arranging the electron pairs (E.P.) in the following manner until all available pairs have been distributed:a) One pair between the central atom and each ligand atom.b) Three more pairs on each outer atom (except hydrogen, which has no additional pairs), yielding 4 E.P. (i.e., an octet) around each ligand atom when the bonding pair is included in the count.c) Remaining electron pairs (if any) on the central atom.5. Calculate the on the central atom.a) Count the electrons shared as bonds. Total = bb) Count the electrons owned as lone pairs.
Total = nc) F = V - (n + b/2), where V = number of valence electrons for the atom.6. If the central atom formal charge is zero or is equal to the charge on the species, the provisional electron distribution from (4) is correct.
Calculate the formal charge of the ligand atoms to complete the Lewis structure.7. If the structure is not correct, calculate the formal charge on each of the ligand atoms. Then to obtain the correct structure, form a multiple bond by sharing an electron pair from the ligand atom that has the most negative formal charge.a) For a central atom from the second (n = 2) row of the periodic table continue this process sequentially until the central atom has 4 E.P. (an octet).b) For all other elements, continue this process sequentially until the formal charge on the central atom is reduced to zero or two double bonds are formed.8. Recalculate the formal charge of each atom to complete the Lewis structure.on to.