Drawing Lewis Structures (A step-by-step Guide by C. Hoeger) ALL …faculty.sdmiramar.edu/choeger/Lewis Structure 231 GA1.pdf · 2008. 4. 14. · Drawing Lewis Structures (A step-by-step

Drawing Lewis Structures (A step-by-step Guide by C. Hoeger)

1 Determine the total number of valence electrons for ALL atoms.

Don’t be concerned with which atom gave what; the SUM TOTAL is what is important. If the entire molecule is charged (i.e. a polyatomic anion or cation) add one valence electron for each unit of negative charge (if it is an anion) and remove one valence electron for each unit of positive charge (if it is a cation).

2 Write a skeleton structure for the molecule, making the least electronegative atom the

central atom.

The order of electronegativity for the nonmetals is F>O>N>Cl>Br>I>S>C>H. Metals will almost always be the central atom if they are present. Hydrogen is never the central atom.

a If there are more than one of the least electronegative atom, your skeletal structure should have those two attached to one another (EXCEPTION: Hydrogen)

3 Connect each member of the skeleton structure to the central atom(s) using a single

line to represent a bond.

Each bond is comprised of two electrons so each line indicates a two-electron bond. 4 Count the number of bonds that you made in step 3 and multiply that by 2. Subtract

that number from the number of total valence electrons from step 1. This is the number of electrons you have left to distribute.

5 Starting with the atoms bonded TO the central atom (the ‘outside’ atoms), distribute

the electrons two at a time until all electrons from step 4 are used OR until each atom has an octet (eight electrons around it); DO NOT give hydrogen any of these ‘left-over’ electrons.

A good rule of thumb is to give EACH non-hydrogen two electrons at a time until

each outside atom has two extras, then give two more to each non-hydrogen until each has four extra electrons and so on. Once all of the ‘outside’ atoms have octets (except for the hydrogens), put remaining electrons on the central atom(s) until all electrons are used or every atom has an octet (or a duet, in the case of H).

AT THIS POINT: Anytime you run out of electrons, look at each atom and determine if it does or does not have an octet. If all atoms have an octet go on to Step 8. If not, proceed to Step 6.

6 If there are too few electrons available to complete octets for all the atoms that need them, make double and/or triple bonds between appropriate atoms.

Remember that in doing this, no atom should lose electrons (i.e. double and triple

bonds SHARE electrons BETWEEN the two involved atoms and as a result neither atom involved in the multiple bond loses electrons, so do not move electrons FROM one atom to another).

a A major exception to this rule is one regarding electrically-neutral Lewis

acids containing Be, Al or B: it is possible to find Be with 4e- around it (BeCl2), and B and Al with 6e

- around them (AlCl3 and BF3, for example) 7 If there are too many electrons available: first, RECOUNT the valence electrons

available (steps 2-4 above); then, after completing octets for all the atoms that need them, place remaining electrons on the central atom IF the central atom is a Period 3 or greater element.

8 After you have completed your Lewis structure, CHECK FOR OCTETS. If all atoms

have an octet, calculate formal charges for all atoms (see below for method).

Remember: if you are drawing the Lewis structure of a cation or an anion, it WILL have a charge (or charges) in it. NOTE: as you get better at drawing Lewis structures, you will develop a ‘feel’ as to whether or not an atom has a formal charge or not. USUALLY, if an atom is not in agreement with the guidelines given below in step 9, it may have a charge and should ALWAYS be checked.

9 Guidelines to follow in writing Lewis Structures (NOTE: there are exceptions to these…) i H follows the duet rule and will have a maximum of 2e- around it and will

therefore only form a single, single bond; ii C, N, O, F always obey the octet rule UNLESS they carry a formal charge; iii O usually forms two covalent bonds (two single or one double);

a) when this is the case, the O should also have 2 unshared pairs of electrons on it (called ‘lone pairs’) and should have no formal charge associated with it;

iv N usually forms three covalent bonds (three single; one double and one single or one triple);

a) when this is the case, the N should also have 1 unshared pair (lone pair) of electrons on it and should have no formal charge associated with it;

v C usually forms four covalent bonds (four single; one double and two single; two double; or one triple and one single BUT NEVER ONE QUADRUPLE BOND!);

a) when this is the case, the C should also have NO unshared pairs of electrons remaining and should have no formal charge associated with it;

vi F always forms only one covalent bond to only one atom at a time and is never the central atom;

a) when this is the case, the F should also have 3 unshared pairs (lone pairs) of electrons on it and should have no formal charge associated with it;

vii Cl, Br and I usually form only one covalent bond to only one atom at a time UNLESS Cl, Br or I is the central atom;

a) when the former is the case (i.e. Cl, Br or I are not the central atom), the Cl, Br or I should each also have 3 unshared pairs of electrons on it and should have no formal charge associated with it;

viii B, Be and Al usually have less than an octet (and as a result they form reactive compounds);

ix Second period elements never exceed the octet; x Third row elements and heavier usually satisfy the octet rule but can expand

their octet to 10 e or 12 e; xi If there are an ODD number of valence electrons TOTAL (from step 1), one

atom in the structure will have an unpaired electron, and that atom will usually be the one that is the LEAST electronegative EXCEPT H (of course…).

FORMAL CHARGE CALCULATION Calculation of formal charge, while not difficult, is the most common mistake made by students at all levels of chemistry when drawing Lewis structures. Formal charge is calculated for each atom independent of any others in the molecule. The procedure is as follows:

1) Determine the number of valence electrons (VE) that the atom of interest has; this is most easily done by using a periodic table;

2) Determine the number of unshared (or non-bonded) electrons (NBE) that the

atom of interest has; 3) Determine the number of bonds (B) that the atom of interest has (remember that

a double bond counts as two and a triple bond counts as three); 4) Formal Charge is then calculated using the following formula:

FC = VE - B - NBE Let’s calculate the formal charge for all the atoms in the Lewis structure for SO2 shown below:

S OO

1 2

For Oxygen 1: FC = 6 - 1 - 6 = -1; For Sulfur: FC = 6 - 3 - 2 = +1;

For Oxygen 2: FC = 6 - 2 - 4 = 0; Thus the complete structure looks like the following:

S OO

1 2

DRAWING ORGANIC MOLECULES

Drawing organic molecules is no different than drawing inorganic molecules: you will still use the rules of Lewis structures, calculate formal charges, etc. The major difference is that you will use some shortcuts for drawing organic molecules that we do not (traditionally) use for inorganic molecules. Before we start, let’s review the rules of valence for neutral atoms that are especially important in organic chemistry;

o Neutral Carbon has 4 valence e– available to make 4 covalent bonds (valence = 4); o Neutral Oxygen has 2 valence e– available to make 2 covalent bonds + 2 lone pairs

(valence = 2); o Neutral Nitrogen has 3 valence e– available to make 3 covalent bonds + 1 lone pairs

(valence = 3); o Neutral Halogen has 1 valence e– available to make 1 covalent bonds + 3 lone pairs

(valence = 1); o Hydrogen has a valence of 1 (one bond);

So in the molecules methane (CH4), ethane (C2H4) and acetylene;

H

C

HH

H

C C

HH

H H

C C HH

the carbons are all neutral, as each has a TOTAL of four bonds. Likewise, in methanol (CH3OH), methyl amine (CH3NH2), acetonitrile (CH3CN) and formaldehyde (CH2O):

H

C

HH

O

H

H

C

HH

N

H

H

H

C

HH

C N

O

C

H H

all the atoms are neutral as well (Note: you would have arrived at the same conclusions had you calculated the formal charge for each atom:

Formal Charge = # valence electrons that atom has – bonds that atom has – unshared electrons that atom has Using the idea of valence is a quick way to determine whether you need to calculate the formal charge or not). Notice that the lone pairs on N and O are not shown. It is understood that they are present (remember that in drawing molecular structures, the lone pairs are not usually shown). With this information in mind we can move on to a discussion of the drawing of organic molecules.

There are three basic ways to draw organic molecules: Extended, or Lewis, structure, Condensed structure and Line-Angle structure. We will examine each of these separately.

EXTENDED STRUCTURE

This is nothing new; in this format EVERY bond to carbon, hydrogen, oxygen, etc. is explicitly shown. It is not necessary to show the geometry at each atom, although you can if you wish. For the molecule butane (C4H10) this can be demonstrated by the structures below:

C

C

C

C

H

H

H

H

H

H

H

HH

H

C C C C

H

H

H

H

H

H

H

H

H

H

A B

The first one (A) shows all atoms and bonds AND the geometry at each atom while the second (B) shows just the connections. Both are acceptable; structure B is usually the one shown (reasons for this will become more evident when you get into Chem 140A!).

Extended structure is fine but requires a lot of writing. Showing all the atoms and the bonds…takes too much time! What about a molecule that has 40 carbons and 80 hydrogens…that’s a lot of ink (or pencil). Is there a quicker way? Yes! CONDENSED STRUCTURE

In condensed format we will show what atoms are attached to what atoms, but we will not explicitly show each bond. For example, the carbon on the end of butane (C4H10) has 3 hydrogens and a carbon attached to it; that carbon has two hydrogens and another carbon attached to it, and so on. As a result it would be easier (and faster) to do the following;

C C C C

H

H

H

H

H

H

H

H

H

H

CH3CH2CH2CH3

same as

Condensed structure is used almost to the exclusion of extended structure (in fact, rarely is the extended Lewis structure ever used unless there is a special reason to do so). It is very easy to use and there are very few pitfalls, but there are some rules to remember:

1. The hydrogens attached to a particular carbon are always written directly next to the carbon to which they are attached (either to the right (usually) or left (less often) of the C);

2. Halogens are treated like hydrogens except they are written after the H’s; 3. If there are groupings of atoms other than hydrogen attached to a carbon they can be shown

using parentheses (and subscript numbers if more than one grouping of the same are present) 4. The hydrogens or other atoms attached to any element other than a carbon are always written

directly to the right of the atom to which they are attached; 5. Any valences left open are used for multiple bond formation.

Lets look at each of these.

For #1: Look at the example given above for butane. Note that it is RARE that H’s are written to the left of a C and this practice is usually ONLY done for the carbon on the FAR left. For #2: Look at the example below:

C C C C

H

H

H

Cl

H

H

H

H

H

Cl

CH3CHClCH2CH2Cl

same as

For #3: See below; note that it is understood that the groups inside the parentheses are attached to the carbon immediately to the left:

C C C C

H

H

H

OH

H

H

OH

H

H

H

CH3CH(OH)CH(OH)CH3

same as

C C C C

H

H

H

C

C

H

H

H

H

Cl

CH3C(CH3)2CH2CH3

same as

H

H

H

H

H H

(CH3)3CCH2CH3OR

Notice that for the second structure there are TWO equivalent ways of representing the structure using condensed structure.

For #4: See the example below, as well as #3 above:

H C C N

H

H

H

H H

H

CH3CH2CH2NH2

same as

NOT CH3CH2CH4N CH3CH2CH2H2NNOR For #5: This is a little harder to see. Let’s use three examples to demonstrate this: formaldehyde, acealdehyde, acetic acid and acetone, respectively.

O

C

H H

O

C

CH3 H

O

C

CH3 OH

CH2O CH3CHO CH3COOH

O

C

CH3 CH3

CH3COCH3

formaldehyde acetaldehyde acetic acid acetone

CH3CO2HOR

Notice that H’s attached TO a C are shown just as before. Let’s start with formaldehyde; notice that the C has three atoms shown attached to it (two H’s and one O). Since carbon normally has a valence of four but only three are shown in formaldehyde, the fourth must be due to a multiple bond to the oxygen. Let’s skip to acetone: if we deal with the C’s and H’s first we get:

CH

HH

CH

HH

CO

O

C

CH3 CH3

acetone

The wavy lines indicate that those carbons are bonded to some other atom. All we have left is one C and one O. The ONLY way to give C four valences and O two is to complete the structure the way we did on the right (in the old days, molecules of this type would be written CH3C(O)CH3, where the (O) indicated a double bond to the C. This representation has all but been abandoned). Carboxylic acids (like acetic acid) are figured out the same way: after taking care of all C’s and H’s, then all H’s on other atoms we get:

CH

HH

O

H

COO

C

CH3 OH

acetic acid

We are once again left with the CO unit, just like in acetone. Condensed structures are nice: fast, convenient, relatively easy, but is there an even faster way? Yes!

LINE-ANGLE STRUCTURE

Line-angle represents the ultimate in simplicity. In line-angle format we do NOT have to EXPLICITLY show ANY C’s or H’s attached to C’s; in fact the only atoms that MUST be shown are any non-carbons AND the H’s that may be attached to non-carbons. In line-angle format we will use an angled line to represent a carbon-carbon bond but will not explicitly show the C’s; since we know that C has a valence of four it will be understood that all valences not shown are filled by H’s. Thus, any time a line ends or two or more lines come together a C must be at that end or intersection. When C is bonded to an element other than C or H, that will also be represented by an angled line but the non-carbon will be explicitly shown. Therefore, the end of a line must represent a C; one valence (bond) is shown therefore it must represent a CH3. When two lines meet we have a C at the intersection of those two lines; since only two bonds are shown it must be a CH2. Let’s look at some examples, using some of the molecules we have already used:

C C C C

H

H

H

H

H

H

H

H

H

H

is

CH3 C CH2 CH2

CH3

CH3

Cl isCl

CH3CH2CH2NH2

NH2is

CH

CH2

CH2

CH2

CH2

CH2

C CH

is

O

C

CH3 OH

O

OH

is

Let’s look in a little more detail at how a correlation is drawn between a line-angle structure and a condensed or extended structure. Using the Cl-containing compound above, it maps out as follows:

CH3 C CH2 CH2

CH3

CH3

Cl

Cl

C with 4 bonds shown: C

C with 2 bonds shown, one to a Clmust be a CH2

End of a line; only one bond shown;must be a CH3

This method of drawing structures is extremely fast once you have practiced it a little bit. Line angle is the preferred method of representing organic molecule structure. REAL-WORLD STRUCTURAL REPRESENTATIONS

What has been laid out for you here are three different, yet equivalent, ways of drawing organic molecules. In reality, ALL THREE representations are used to some degree or another. A molecule may, for the most part, be drawn in line angle but have portions that are done in condensed and/or extended format. Some examples are shown below:

BrH

C(CH3)3H

C

C

CH3

H

H

O

NH2

CH3

H

Notice that in all of the structures shown there is SOME of each motif. Now, why would we mix our formats? To answer that question, look quickly at the structures above. What is the first thing that catches your eye? Odds are it is the portion of the molecule drawn in either extended or condensed format. This is one of the reasons for mixed format: to draw your attention to a part of the molecule that bears some importance in a discussion that follows. Other reasons exist but this is the most important for purposes of this discussion. Whatever format or mixed format you like is your own business, but you do need to be able to move back and forth between the three formats on a regular basis.

Chemistry 231L Group Activity # 1 – 15 points (Please turn in Last 3 pages) Instructions for Part I: Lewis Dot Structure, Formal Charge, and VSEPR Hints on drawing Lewis Dot Structure:

1. See attached “Drawing Lewis Structures” and “Drawing Organic Molecules” Instructions for Part II: Drawing Organic Molecules

A. Molecular Formula a. It only offers molecular composition information and nothing else. It will be up to you to determine

connectivity of molecule b. Example: C4H6O3; C’s are almost always attached to C’s BUT it depends on functional groups present.

B. Kekulé Structure or Line bond (Lewis extended) Structure a. This is what you will usually draw when asked to draw Lewis Dot Structure, it would be ok to omit the

lone pairs when asked to draw Kekulé structure, but it will not be correct if lone pairs are not drawn in for Lewis Dot Structure. It is very helpful to keep drawing the lone pairs. Lone pairs are often not drawn BUT understood.

b. This representation also does not offer any 3-Dimensional information, but it does offer some molecular geometry information.

c. Example:

H

C

C

O

C

C

H

HH

HH

O O

H

C

C

O

C

C

H

HH

HH

O O

C. Wedged-Line Notation: a. Shows the 3-D structure of the molecule, with lines = bonds in plane, dash = bonds pointing into the

plane wedge = bond coming out of the plan b. Example: (lone pairs could be omitted)

H

C

C

O

C

C

H

H HHH

O O D. Condensed Structure

a. Offers basic information about connectivity b. Example: CH3COOC(O)CH3, (O) is =O; (OH) is O–H; CHO is C(O)H

E. Skeletal or Line-Angle Structure a. This representation is the simplest but it has rules that you must follow. Text p 55 and see attached

“Drawing Organic Molecules” 1. Carbon atoms aren’t usually shown, Instead, a carbon atom is assumed to be at each

intersection of two lines (bonds) and at the end of each line. Occasionally, a carbon atom might be indicated for emphasis or clarity.

2. Hydrogen atoms bonded to carbons aren’t shown. Since carbon always has a valence of 4, we mentally supply the correct number of hydrogen atoms for each carbon.

3. Atoms other than carbon and hydrogen are shown. b. Example

O

O O

This exercise is to get you to start to become fluent at drawing Lewis structures. You will be learning more in chapter three. For now you will draw structures using the four representations above to gain more practice.

2

Name: ______________________________________ Name: ________________________________ Part I. Lewis Dot Structure, Formal Charge, and VSEPR

PRACTICE

1) Draw Lewis Dot Structures in VSEPR electronic geometry format from the following condensed formula; use dash and wedge where it is necessary to show 3D.

2) Specify polar or nonpolar BOND with “P” or “NP” 3) Specify polar or nonpolar MOLECULE, if molecule is polar show approx. direction of polarity 4) Show Formal Charge where ever it is appropriate

NI3 H2O

CH3OH Cl2

*CH3COO– BF3NH3

CO2 *CHClCClCH3 (put Cl’s on same side)

*=Link Carbons in skeleton

3

Part II. Drawing Organic Molecules

Condensed Formula Extended Structure Line-Angle

CH3CH2CH3

CH3OC(CH3)3

HCCCH2OH

CH3C

O

CHCHOCH3

HOCH2CH2OCH2CH2OH

For last molecule, draw a wedge-dash structure on the back of this page:

4

Part III. Hybridization

1. Allene, CH2CCH2. Note that all C’s are connected a. Draw the Lewis Dot structure of allene, CH2CCH2 in the box on the left b. Label the carbons with the numbers 1-3 c. Which carbon(s) is(are) sp hybridized according to your labeling? __________

d. Which carbon(s) is(are) sp2 hybridized according to your labeling?__________

e. What is the bond angle for H–C–H? ________

f. What is the bond angle for C–C–C?_________

g. What is the bond angle for H–C–C?_________

h. Re-draw below the stucture you drew for (a). Label each bond as either σ or π. Show all your bond angles

i. Show an orbital picture of the π bond framework. Be sure to label the orbitals properly and the bond angle is reflected in your drawing.

j. Re-draw allene in a wedged-dash notation in the box provided above (HINT: Build a model first!)

2. a) What kind of hybridization would you expect for the following N atoms in these two molecules?

NH2

aniline

N

pyridine

B) For which of these molecules are the nitrogen’s lone pairs LEAST available for ? Why?

a. j.