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Drawing Structures of Organic Molecules...

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Part 01: Drawing Structures of Organic Molecules I.

Bond Formation: The Octet Rule

•

Atoms transfer or share electrons to have the same number of electrons as the nearest noble gas atom.

•

Noble gas atoms have a completely filled valence shell of electrons.

•

Most groups will end up having 8 valence electrons to completely fill the valence shell. This is the octet rule.

•

Exceptions include hydrogen, where it only needs 2 electrons, boron and aluminum are stable with only 6 valence electrons, and elements beyond the third row can have an expanded octet of more than 8 electrons. Ionic bonds (transfer electrons)

•

Sometimes atoms satisfy the octet rule by transferring electrons from one atom to another.

Li

Cl

Li

Cl Opposite charges attract

•

This forms ions with opposite charges, resulting in attraction and ionic bond formation.

When do we form ionic bonds? •

The more electronegative an atom is, the more it is able to attract electrons.

•

Electronegativity increases from left to right and from bottom to top of the periodic table. Electronegativity Increases

•

Ionic bonds form when the difference in electronegativity is larger than 2.0 (between metals and non-metals).

Covalent bonds (share electrons) •

Another strategy is to share valence electrons between atoms.

Cl

Cl

Cl Cl Electrons are shared

•

The two atoms then come together forming a covalent bond.

•

This type of bonding is commonly seen with atoms closer in electronegativity.

•

Let's take a look at how C, N, O, H, and halogens form covalent bonds.

Atom

Valence Electrons

To Complete Octet

Carbon (C)

4

4 e- needed

Nitrogen (N)

5

3 e- needed

Oxygen (O)

6

2 e- needed

Hydrogen (H)

1

1 e- needed

Halogen

7

1 e- needed

Number of Bonds in Neutral compounds (number of shared electron pairs) C

C

C

N

N

N

4

3

2

Easy way to remember HONC 1234 rule! Hydrogen and Halogen make 1 bond Oxygen make 2 bonds Nitrogen make 3 bonds Carbon make 4 bonds

O

1

O

H

Cl

1

For charged molecules, the number of bonds change:

Atom

Number of Bonds to Neutral Atom

Carbon (C)

4

Number of Bonds when Cationic

C

3

Number of Bonds when Anionic

C

C

Carbocation (they are special!)

Nitrogen (N)

3

Oxygen (O)

2

N

N

O

O

O

2

N

1

3

N

O

Let's break this down further: Neutral C atom: 4 electrons in valence shell C+: __ electrons C–: __ electrons

*can bond to __ atoms *can bond to __ atoms (similar to neutral N atom)

Neutral O atom: 6 electrons in the valence shell O+: __ electrons O–: __electrons Q. What about Nitrogen? A.

C

Carbanion

N

4

C

3

*can bond to __ atoms (similar to C– and N) *can bond to __ atoms

A shortcut to calculate formal charge: Formal Charge = [# of electrons in neutral atom] – [# of electrons an atom has]

H H H C N O H H We are only counting 1 electron in each bond Let’s do some exercises!

H O C

N

What is the formal charge for carbon? What is the formal charge for oxygen? What is the formal charge for nitrogen? (A) +1 (B) 0 (C) -1

Let’s use Plickers! • • •

Show me your answer by holding up the card with the correct answer at the top Do not cover up the square when holding the card up Everybody has a different card!

Exceptions to the octet rule: 1. Third row and higher elements. These have d orbitals available for binding, allowing the atoms to exceed an octet in their valence shell. Examples: Sulfuric Acid (H2SO4) and Phosphoric Acid (H3PO4)

O H O S O H O

O H O P O H O H

Sulfur and phosphorus both have expanded octets.

2. Atoms with open shells (Incomplete octets). In general, atoms in Group III, such as boron and aluminum, will have an open shell.

F B Can make 3 shared bonds

only ____ electrons around boron

Open shell compounds can accept a pair of electrons to complete an octet. F F B F

II.

H N H H

Lewis structures

Shows connectivity and the location of all bonding and non-bonding electrons and formal charges. But do not show actual 3D arrangement of structure.

To draw Lewis structure: • • •

Follow the HONC-1234 rule for neutral compounds. Maximize the number of bonds without breaking the octet rule. Assign formal charges to any charged atoms.

Example 1: draw the Lewis structure for CH3Br. First Method: GChem way Calculate total number of valence electrons. Arrange pairs of electrons around each atom to achieve octet. If any atom does not have octet, add a double bond using a lone pair. C: __ e3H: __ eBr: __e-

draw all nonbonding electrons on Lewis structures! don't use lines for lone pairs in OChem

Second Method: Shortcut (Ochem way) C: __ bonds H: __ bond Br:__ bond Example 2: Draw the Lewis structure for ethylene (CH2CH2). GChem method: 2C: __ e4H: __ eShortcut method: How many bonds does a neutral carbon form? C: __ bonds H: __ bond Which method do you prefer? (A) G Chem method (B) O Chem method Example 3: Draw the Lewis structure for NaOCH2CH3 This compound has a combination of covalent and ionic bonds

Example 4: Write the Lewis structure for CH3CHOH+ 2C: __ e5 H: __ eO: __eRemember to assign charges to atoms! Formal Charge = [# of electrons in neutral atom] – [# of electrons an atom has] Important points about this example: 1. Assign formal charge to any atom that has a formal charge. Formal Charge = [# of electrons in neutral atom] – [# of electrons an atom has] 2. The above molecule can be represented by more than one Lewis structure. They are resonance structures for the same molecule.

H H H C C O H H

H H H C C O H H

• resonance structures differ only in the arrangement of electrons. NOT in the connectivity. • the actual structure of the molecule is a resonance hybrid of the two structures. Arrow review: resonance arrow equilibrium arrow reaction arrow curvy arrow — shows movement of a pair of electrons fishhook arrow — shows movement of a single electron One way we can draw the resonance hybrid:

H H H C C O H H fill in partial charges and bonds in the structure above

Important points about resonance structures: •

Resonance structures are not real. An individual resonance structure does not accurately represent the structure of the molecule or ion. Only the hybrid does.

•

Resonance structures are not in equilibrium with each other. There is no movement of electrons from one form to another.

•

Resonance structures are not isomers. Two isomers differ in the arrangement of both atoms and electrons, whereas resonance structures differ only in the arrangement of electrons.

H H C C O H

and

H H C C O H

H O H H C C H

and

H O H C C H H

Electron pair moves to a different location

Atom is moving to a different location

Therefore: _______________________

Therefore: _______________________ (atoms are bonded together differently but same molecular formula)

3. These two resonance structures are not equivalent:

H H H H C C O H

_______ resonance structure

H H H H C C O H

_______ resonance structure

__ covalent bonds

__covalent bonds

______________________ octet

______________________ octet

(+) charge on ______ electronegative atom

(+) charge on _______ electronegative atom

Q: What makes a good resonance structure? A: The best resonance structures are the most stable!

Look for the following structural features: IN ORDER OF IMPORTANCE: Rule 1 — resonance structures with more bonds and fewer charges are more stable. Rule 2 — resonance structures in which every atom has an octet are more stable. Rule 3 — resonance structures that place a negative charge on a more electronegative atom are more stable. ADDITIONAL RULES *only CARBON without octet! (never have oxygen, nitrogen, halogens without octet) *NEVER 2+ or 2- charge on any atom! III.

Shorthand Representations

A. Partially Condensed Structures In partially condensed structures, the bonds between carbon and hydrogen are not drawn out. Instead, the hydrogens attached to a given carbon are listed after that carbon in the structure.

H H H H H C C C C Cl H H H H

• •

=

H O H C C H H

=

H O H H C C C H H H

=

H O H C C O H H

=

The bonds between carbons and to other atoms are still explicitly shown. This makes it easy to see the structure of the carbon skeleton and any branches. This is a convenient way to save time and space when writing structures, and we will use it most of the time.

B. Fully Condensed Structures In fully condensed structures, all bond dashes are omitted. All the atoms or groups attached to a carbon are listed immediately after that carbon. Complicated, polyatomic groups and branches are written in parentheses to show that they are together.

H H H H H C C C C Cl H H H H

•

•

=

H O H C C H H

=

H O H H C C C H H H

=

H O H C C O H H

=

NOT:

NOT:

NOT:

Fully condensed structures save a great deal of space, and they are easy to type out. However, they are most useful only for relatively simple structures, and they are very confusing when the carbon chain has complicated branches. This type of structure will only be used for simple molecules with few branches.

C. 3-D Structures 3-D gives the three dimensional orientation space. Example:

D. Bond-Line (Skeletal) Structures 1. In bond-line structures, all of the carbon and hydrogen atoms are omitted. 2. Only the bonds between carbons are drawn. 3. All the vertices represent carbon atoms, and they are assumed to be bonded to enough hydrogens to complete their octet. 4. Also, the end of any line that is not attached to a written out atom is also a carbon.

H H H H H C C C C Cl H H H H

•

=

H O H C C H H

=

H O H H C C C H H H

=

H O H C C O H H

=

Bond-line structures focus attention on the shape of the carbon skeleton, and they also save time and space. They are used especially for compounds that contain rings or complicated carbon skeletons.

It is common to combine more than one type of structure:

VI.

Functional Groups of Organic Molecules

A functional group is an atom or group of atoms with characteristic chemical and physical properties. A functional group contains either heteroatoms or p-bonds (double bonds). Heteroatom: An atom other than carbon or hydrogen. Common heteroatoms are nitrogen, oxygen, sulfur, phosphorus, and halogens. p -bonds: the most common p-bonds occur in C–C and C–O double bonds, and C–C and C–N triple bonds.

How do heteroatoms and p-bonds confer reactivity on a molecule? Heteroatoms have lone pairs and create electron deficient or electron rich sites on a molecule. p-bonds are easily broken in chemical reactions. C – C double bonds create electron rich sites in a molecule. A C – C double bond makes a molecule a base and a nucleophile. Lone pairs make oxygen electron rich

p-bond makes this compound electron rich

C O

C C

Electron deficient carbon C–C and C– H s-bonds (single bond) in organic molecules form the carbon backbone to which functional groups are bonded. C–C and C–H s-bonds are typically not reactive, and as a result, they are not considered functional groups. They are often designated as R, especially when the focus is on the functional group. R — functional group carbon skeleton (unreactive portion of molecule)

reactive portion of molecule...