Drawing Lewis Structures (A Step-by-step Guide By C .
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 isimportant. If the entire molecule is charged (i.e. a polyatomic anion or cation) addone valence electron for each unit of negative charge (if it is an anion) and removeone 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 thecentral 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 isnever the central atom.a If there are more than one of the least electronegative atom, your skeletalstructure 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 singleline 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. Subtractthat number from the number of total valence electrons from step 1. This is thenumber of electrons you have left to distribute.5 Starting with the atoms bonded TO the central atom (the ‘outside’ atoms), distributethe electrons two at a time until all electrons from step 4 are used OR until each atomhas an octet (eight electrons around it); DO NOT give hydrogen any of these ‘leftover’ electrons.A good rule of thumb is to give EACH non-hydrogen two electrons at a time untileach outside atom has two extras, then give two more to each non-hydrogen untileach 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 allelectrons 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 determineif 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 needthem, make double and/or triple bonds between appropriate atoms.Remember that in doing this, no atom should lose electrons (i.e. double and triplebonds SHARE electrons BETWEEN the two involved atoms and as a result neitheratom involved in the multiple bond loses electrons, so do not move electronsFROM one atom to another).a A major exception to this rule is one regarding electrically-neutral Lewisacids 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 electronsavailable (steps 2-4 above); then, after completing octets for all the atoms that needthem, place remaining electrons on the central atom IF the central atom is a Period 3or greater element.8 After you have completed your Lewis structure, CHECK FOR OCTETS. If all atomshave 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 WILLhave a charge (or charges) in it. NOTE: as you get better at drawing Lewisstructures, you will develop a ‘feel’ as to whether or not an atom has a formalcharge or not. USUALLY, if an atom is not in agreement with the guidelinesgiven 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 tothese )iH follows the duet rule and will have a maximum of 2e- around it and williiiiiivvviviiviiiixxxitherefore only form a single, single bond;C, N, O, F always obey the octet rule UNLESS they carry a formal charge;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 ofelectrons on it (called ‘lone pairs’) and should have no formal chargeassociated with it;N usually forms three covalent bonds (three single; one double and one singleor one triple);a) when this is the case, the N should also have 1 unshared pair (lonepair) of electrons on it and should have no formal charge associated withit;C usually forms four covalent bonds (four single; one double and two single;two double; or one triple and one single BUT NEVER ONE QUADRUPLEBOND!);a) when this is the case, the C should also have NO unshared pairs ofelectrons remaining and should have no formal charge associated with it;F always forms only one covalent bond to only one atom at a time and is neverthe central atom;a) when this is the case, the F should also have 3 unshared pairs (lonepairs) of electrons on it and should have no formal charge associated withit;Cl, Br and I usually form only one covalent bond to only one atom at a timeUNLESS 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 itand should have no formal charge associated with it;B, Be and Al usually have less than an octet (and as a result they form reactivecompounds);Second period elements never exceed the octet;Third row elements and heavier usually satisfy the octet rule but can expandtheir octet to 10 e or 12 e;If there are an ODD number of valence electrons TOTAL (from step 1), oneatom in the structure will have an unpaired electron, and that atom will usuallybe the one that is the LEAST electronegative EXCEPT H (of course ).
FORMAL CHARGE CALCULATIONCalculation of formal charge, while not difficult, is the most common mistake made bystudents at all levels of chemistry when drawing Lewis structures. Formal charge iscalculated for each atom independent of any others in the molecule. The procedure is asfollows: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 theatom of interest has;3) Determine the number of bonds (B) that the atom of interest has (remember thata 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 - NBELet’s calculate the formal charge for all the atoms in the Lewis structure for SO2 shownbelow:OS1O2For 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:O1SO2
DRAWING ORGANIC MOLECULESDrawing organic molecules is no different than drawing inorganic molecules: you will still use therules of Lewis structures, calculate formal charges, etc. The major difference is that you will use someshortcuts 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 inorganic 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;HHHCCHHHCHHCHCHthe carbons are all neutral, as each has a TOTAL of four bonds. Likewise, in methanol (CH3OH),methyl amine (CH3NH2), acetonitrile (CH3CN) and formaldehyde (CH2O):HHHCHHOHHCHHNHOCHCNCHHHall the atoms are neutral as well (Note: you would have arrived at the same conclusions had youcalculated the formal charge for each atom:Formal Charge # valence electrons that atom has – bonds that atom has – unshared electrons that atom hasUsing the idea of valence is a quick way to determine whether you need to calculate the formal chargeor 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 thisinformation 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, Condensedstructure and Line-Angle structure. We will examine each of these separately.
EXTENDED STRUCTUREThis is nothing new; in this format EVERY bond to carbon, hydrogen, oxygen, etc. is explicitlyshown. It is not necessary to show the geometry at each atom, although you can if you wish. For themolecule butane (C4H10) this can be demonstrated by the structures below:HHHHCCCCHHHHHHHHHHHCCCCHHHHHBAThe 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 thiswill 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 takestoo much time! What about a molecule that has 40 carbons and 80 hydrogens that’s a lot of ink (orpencil). Is there a quicker way? Yes!CONDENSED STRUCTUREIn condensed format we will show what atoms are attached to what atoms, but we will not explicitlyshow each bond. For example, the carbon on the end of butane (C4H10) has 3 hydrogens and a carbonattached to it; that carbon has two hydrogens and another carbon attached to it, and so on. As a resultit would be easier (and faster) to do the following;HHHHHCCCCHHHHsame asCH3CH2CH2CH3H
Condensed structure is used almost to the exclusion of extended structure (in fact, rarely is theextended Lewis structure ever used unless there is a special reason to do so). It is very easy to use andthere 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 towhich 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 shownusing 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 writtendirectly 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 theleft of a C and this practice is usually ONLY done for the carbon on the FAR left.For #2: Look at the example below:HHClHHCCCCHHHHClsame asCH3CHClCH2CH2ClFor #3: See below; note that it is understood that the groups inside the parentheses are attached to thecarbon immediately to the left:HHHHOHHHCCCCHHOHHHCHHHCCHHCHHCCHHClHHsame asCH3CH(OH)CH(OH)CH3same asCH3C(CH3)2CH2CH3OR(CH3)3CCH2CH3Notice that for the second structure there are TWO equivalent ways of representing the structureusing condensed structure.
For #4: See the example below, as well as #3 above:HHHCCNHHHHsame asCH3CH2CH2NH2CH3CH2CH4NNOTNORCH3CH2CH2H2NFor #5: This is a little harder to see. Let’s use three examples to demonstrate this: formaldehyde,acealdehyde, acetic acid and acetone, 3COOH OR CH3CO2HCH3COCH3formaldehydeacetaldehydeacetic acidacetoneNotice that H’s attached TO a C are shown just as before. Let’s start with formaldehyde; notice thatthe C has three atoms shown attached to it (two H’s and one O). Since carbon normally has a valenceof four but only three are shown in formaldehyde, the fourth must be due to a multiple bond to theoxygen. Let’s skip to acetone: if we deal with the C’s and H’s first we get:OCOHCHHCHCCH3HHCH3acetoneThe wavy lines indicate that those carbons are bonded to some other atom. All we have left is oneC and one O. The ONLY way to give C four valences and O two is to complete the structure the waywe 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). Carboxylicacids (like acetic acid) are figured out the same way: after taking care of all C’s and H’s, then all H’son other atoms we get:OCOHHCHCOHCH3OHacetic acidWe 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 STRUCTURELine-angle represents the ultimate in simplicity. In line-angle format we do NOT have toEXPLICITLY show ANY C’s or H’s attached to C’s; in fact the only atoms that MUST be shown areany non-carbons AND the H’s that may be attached to non-carbons. In line-angle format we will usean angled line to represent a carbon-carbon bond but will not explicitly show the C’s; since we knowthat 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. WhenC is bonded to an element other than C or H, that will also be represented by an angled line but thenon-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 theintersection of those two lines; since only two bonds are shown it must be a CH2. Let’s look at someexamples, using some of the molecules we have already used:HHHHHCCCCHHHHHisCH2 CH2 ClisCH3CH2CH2NH2isCH3CH3 CClCH3CH2CH2CH2CH2CNH2CHCHisCH2OOCCH3OHisOH
Let’s look in a little more detail at how a correlation is drawn between a line-angle structure and acondensed or extended structure. Using the Cl-containing compound above, it maps out as follows:C with 2 bonds shown, one to a Clmust be a CH2End of a line; only one bond shown;must be a CH3ClC with 4 bonds shown: CCH3CH3 CCH2 CH2 ClCH3This method of drawing structures is extremely fast once you have practiced it a little bit. Lineangle is the preferred method of representing organic molecule structure.REAL-WORLD STRUCTURAL REPRESENTATIONSWhat has been laid out for you here are three different, yet equivalent, ways of drawing organicmolecules. In reality, ALL THREE representations are used to some degree or another. A moleculemay, for the most part, be drawn in line angle but have portions that are done in condensed and/orextended format. Some examples are shown below:BrHOHHCCHCH3CH3HC(CH3)3NH2Notice that in all of the structures shown there is SOME of each motif. Now, why would we mixour formats? To answer that question, look quickly at the structures above. What is the first thing thatcatches your eye? Odds are it is the portion of the molecule drawn in either extended or condensedformat. This is one of the reasons for mixed format: to draw your attention to a part of the moleculethat bears some importance in a discussion that follows. Other reasons exist but this is the mostimportant for purposes of this discussion. Whatever format or mixed format you like is your ownbusiness, but you do need to be able to move back and forth between the three formats on a regularbasis.
Chemistry 231L Group Activity # 1 – 15 points (Please turn in Last 3 pages)Instructions for Part I: Lewis Dot Structure, Formal Charge, and VSEPRHints on drawing Lewis Dot Structure:1. See attached “Drawing Lewis Structures” and “Drawing Organic Molecules”Instructions for Part II: Drawing Organic MoleculesA. Molecular Formulaa. It only offers molecular composition information and nothing else. It will be up to you to determineconnectivity of moleculeb. 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) Structurea. This is what you will usually draw when asked to draw Lewis Dot Structure, it would be ok to omit thelone pairs when asked to draw Kekulé structure, but it will not be correct if lone pairs are not drawn in forLewis Dot Structure. It is very helpful to keep drawing the lone pairs. Lone pairs are often not drawnBUT understood.b. This representation also does not offer any 3-Dimensional information, but it does offer some moleculargeometry information.c. Example:HHHCHOHHCCCOOHHCHHOHCCCOOHC. Wedged-Line Notation:a. Shows the 3-D structure of the molecule, with lines bonds in plane, dash bonds pointing into theplane wedge bond coming out of the planb. Example: (lone pairs could be omitted)HHHCHOHCCCOOHD. Condensed Structurea. Offers basic information about connectivityb. Example: CH3COOC(O)CH3, (O) is O; (OH) is O–H; CHO is C(O)HE. Skeletal or Line-Angle Structurea. 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 eachintersection of two lines (bonds) and at the end of each line. Occasionally, a carbon atommight be indicated for emphasis or clarity.2. Hydrogen atoms bonded to carbons aren’t shown. Since carbon always has a valence of 4, wementally supply the correct number of hydrogen atoms for each carbon.3. Atoms other than carbon and hydrogen are shown.b. ExampleOOOThis exercise is to get you to start to become fluent at drawing Lewis structures. You will be learning more in chapterthree. For now you will draw structures using the four representations above to gain more practice.
2Name:Name:Part I. Lewis Dot Structure, Formal Charge, and VSEPRPRACTICE1) Draw Lewis Dot Structures in VSEPR electronic geometry format from the following condensed formula; usedash 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 polarity4) Show Formal Charge where ever it is ClCH3 (put Cl’s on same side)* Link Carbons in skeleton
3Part II. Drawing Organic MoleculesCondensed FormulaExtended StructureCH3CH2 CH3CH3OC(CH3)3HCCCH2OHOCH3 C CHCHOCH3HOCH2CH2OCH2CH2OHFor last molecule, draw a wedge-dash structure on the back of this page:Line-Angle
4Part III. Hybridization1. Allene, CH2CCH2. Note that all C’s are connecteda.Draw the Lewis Dot structure of allene, CH2CCH2 in the box on the lefta.j.b.c.Label the carbons with the numbers 1-3Which 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 anglesi.Show an orbital picture of the π bond framework. Be sure to label the orbitals properly and the bondangle 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?NH2NanilinepyridineB) For which of these molecules are the nitrogen’s lone pairs LEAST available for ? Why?
CN) and formaldehyde (CH 2 O): H C H H O H H C H H N H H H C C N O C H H . 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 Lewis structure ever used u
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Atomic Theory Unit – Drawing Lewis Structures I. Chemists follow several guidelines when visually representing covalent compounds. The conventional method of drawing produces Lewis structures. Work together to draw the compounds using these Lewis struc
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