Organic Chemistry - Chapter 1 Lecture Notes

Atomic Structure

• The nucleus contains positively charged protons and uncharged neutrons.
• The electron cloud is composed of negatively charged electrons.
• The atomic number is the number of protons in the nucleus and also the number of electrons surrounding (i.e., protons = electrons).
• The mass number is the number of protons plus neutrons in the nucleus.
  • Example: Carbon has six protons and six neutrons. Carbon’s atomic number is 6; its mass number is 12.

Ions

• A cation is positively charged and has fewer electrons than protons.
• An anion is negatively charged and has more electrons than protons.

Isotopes

• Isotopes are two atoms of the same element having a different number of neutrons.
• Isotopes have different mass numbers.
• Most isotopes of carbon have 6 protons and 6 neutrons.
• The atomic weight of a particular element is the weighted average of the mass of all its isotopes reported in atomic mass unit (amu).

Periodic Table

• Elements in the same row are similar in size.
• Elements in the same column have similar electronic and chemical properties.
• First Row
  • There is only one orbital in the first shell.
  • Each shell can hold a maximum of two electrons.
  • Therefore, there are two elements in the first row: H and He.
• Second Row
  • Each element in the second row of the periodic table also has four orbitals available to accept additional electrons: one 2s orbital, and three 2p orbitals.
  • Each of the four orbitals in the second shell hold two electrons.
  • There is a maximum capacity of eight valence electrons for elements in the second row.
  • The second row of the periodic table consists of eight elements, obtained by adding electrons to the 2s and three 2p orbitals.

Atomic Orbitals

• An s orbital has a sphere of electron density and is lower in energy than the other orbitals of the same shell.
• A p orbital has a dumbbell shape and contains a node (no electron density) at the nucleus. It is higher in energy than an s orbital.

Bonding

• Bonding is the joining of two atoms in a stable arrangement.
• Through bonding, atoms gain, lose, or share electrons to attain the electronic configuration of the noble gas closest to them in the periodic table.
• Atoms can form either ionic or covalent bonds to attain a complete outer shell (octet rule for second row elements).
  • Ionic bonds result from the transfer of electrons from one element to another.
  • Covalent bonds result from the sharing of electrons between two nuclei.

Ionic Bonding

• An ionic bond generally occurs when elements on the far left side of the periodic table combine with elements on the far right side, ignoring noble gases.
• A positively charged cation formed from the element on the left side attracts a negatively charged anion formed from the element on the right side (e.g., sodium chloride, NaCl).

Covalent Bonding

• Covalent bonding occurs with elements like carbon in the middle of the table with elements that have similar electronegativity.
• Covalent bonds also occur between two of the same elements from the sides of the table.
• A covalent bond is a two-electron bond, and a compound with covalent bonds is called a molecule.
• Hydrogen forms one covalent bond.
• When two hydrogen atoms are joined in a bond, each has a filled valence shell of two electrons.
• Second row elements can have no more than eight electrons around them.
  • For neutral molecules, this has two consequences:
    • Atoms with one, two, three, or four valence electrons form one, two, three, or four bonds, respectively, in neutral molecules (e.g., $BF<em>3$, $CH</em>4$).
    • Atoms with five or more valence electrons form enough bonds to give an octet (e.g., $NH_3$). In this case, predicted number of bonds = 8 − number of valence electrons.

Nonbonded Electrons

• When second-row elements form fewer than four bonds their octets consist of both bonding (shared) and nonbonding (unshared) electrons.
• Unshared electrons are also called lone pairs.

Lewis Structures

• Lewis structures are electron dot representations for molecules.

General rules for drawing Lewis structures:

• Draw only the valence electrons.
• Give every second-row element no more than eight electrons.
• Give each hydrogen two electrons.
• In a Lewis Structure, a solid line indicates a two- electron covalent bond.

How to Draw a Lewis Structure

• Step [1] Arrange atoms next to each other that you think are bonded together.
  • Always place hydrogen and halogens on the periphery because they form only one bond each.
  • Place no more atoms around an atom than the number of bonds it usually forms.
• Step [2] Count the electrons.
  • Count the number of valence electrons from all atoms.
  • Add one electron for each negative charge.
  • Subtract one electron for each positive charge.
  • This gives the total number of electrons that must be used in drawing the Lewis structure.
• Step [3] Arrange the electrons around the atoms.
  • Place a bond between every two atoms, giving two electrons to each H and no more than eight to any second-row atom.
  • Use all remaining electrons to fill octets with lone pairs.
  • If all valence electrons are used and an atoms does not have an octet, form multiple bonds.
• Step [4] Assign formal charges to all atoms.

Multiple Bonds

• If all valence electrons are used and an atom does not have an octet, form multiple bonds.
• To give both C’s an octet, change one lone pair into one bonding pair between the two C’s, forming a double bond.

Formal Charge

• Formal charge is the charge assigned to individual atoms in a Lewis structure.
• Formal charge is calculated as follows: The number of electrons “owned” by an atom is determined by its number of bonds and lone pairs.
• An atom “owns” all of its unshared electrons and half of its shared electrons.

Isomers

• Sometimes more than one arrangement of atoms (Lewis structure) is possible for a given molecular formula.
• These two compounds are called isomers.
• Isomers are different molecules having the same molecular formula.
• Ethanol and dimethyl ether are constitutional isomers.

Exceptions to the Octet Rule

• Elements in Groups 2A and 3A
• Elements in the Third Row

Resonance

• Some molecules cannot be adequately represented by a single Lewis structure.
• These structures are called resonance structures or resonance forms.
• A double-headed arrow is used to separate the two resonance structures.
• Resonance structures are two Lewis structures having the same placement of atoms but a different arrangement of electrons.
• Neither resonance structure is an accurate representation for (HCONH)− .
• The true structure is a composite of both resonance forms and is called a resonance hybrid.
• The hybrid shows characteristics of both structures.
• Resonance allows certain electron pairs to be delocalized over two or more atoms, and this delocalization adds stability.
• A molecule with two or more resonance forms is said to be resonance stabilized.

Basic Principles of Resonance Theory

• Resonance structures are not real. An individual resonance structure does not accurately represent the structure of a 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.

Drawing Resonance Structures

• Rule [1]: Two resonance structures differ in the position of multiple bonds and nonbonded electrons. The placement of atoms and single bonds always stays the same.
• Rule [2]: Two resonance structures must have the same number of unpaired electrons.
• Rule [3]: Resonance structures must be valid Lewis structures. Hydrogen must have two electrons and no second-row element can have more than eight electrons.

Curved Arrow Notation

• Curved arrow notation is a convention that shows how electron position differs between two resonance forms.
• Curved arrow notation shows the movement of an electron pair.
• The tail of the arrow always begins at the electron pair, either in a bond or lone pair.
• The head points to where the electron pair “moves.”

Atoms Without Octets

• Resonance structures can have an atom with fewer than 8 electrons.
• However, resonance structures can never have a second- row element with more than 8 electrons.
• Two different resonance structures can be drawn when a lone pair is located on an atom directly bonded to a double bond.

The Resonance Hybrid

• A resonance hybrid is a composite of all possible resonance structures.
• In the resonance hybrid, the electron pairs drawn in different locations in individual resonance forms are delocalized.
• When two resonance structures are different, the hybrid looks more like the “better” resonance structure.
• The “better” resonance structure is called the major contributor to the hybrid, and all others are minor contributors.
• A “better” resonance structure is one that has more bonds and fewer charges.

Molecular Shape

• Two variables define a molecule’s structure: bond length and bond angle.
• Bond length decreases across a row of the periodic table as the size of the atom decreases.
• Bond length increases down a column of the periodic table as the size of an atom increases.

Molecular Geometry

• The number of groups surrounding a particular atom determines its geometry. A group is either an atom or a lone pair of electrons.
• The most stable arrangement keeps these groups as far away from each other as possible. This is exemplified by Valence Shell Electron Pair Repulsion (VSEPR) theory.
  • Two groups: linear, 180° bond angle
  • Three groups: trigonal planar, 120° bond angle
  • Four groups: tetrahedral, 109.5° bond angle
• A solid line is used for a bond in the plane.
• A wedge is used for a bond in front of the plane.
• A dashed line is used for a bond behind the plane
• Note that wedges and dashed wedges are used for groups that are really aligned one behind another.
• A Nonbonded Pair of Electrons is Counted as a “Group”
  • In ammonia (NH3), one of the four groups attached to the central N atom is a lone pair. The group geometry is a tetrahedron. The molecular shape is referred to as trigonal pyramidal.
  • In water (H2O), two of the four groups attached to the central O atom are lone pairs. The group geometry is a tetrahedron. The molecular shape is referred to as bent.
• In both NH3 and H2O the bond angle is smaller than the theoretical tetrahedral bond angle because of repulsion of the lone pairs of electrons. The bonded atoms are compressed into a smaller space with a smaller bond angle.

Drawing Organic Molecules

*Condensed Structures

• All atoms are drawn in, but the two-electron bond lines are generally omitted.
• Atoms are usually drawn next to the atoms to which they are bonded.
• Parentheses are used around similar groups bonded to the same atom.
• Lone pairs are omitted.

Skeletal Structures

• Assume there is a carbon atom at the junction of any two lines or at the end of any line.
• Assume there are enough hydrogens around each carbon to make it tetravalent.
• Draw in all heteroatoms and the hydrogens directly bonded to them.
• A charge on a carbon atom takes the place of one hydrogen atom. The charge determines the number of lone pairs; negatively charged carbon atoms have one lone pair and positively charged carbon atoms have none.
• Skeletal structures often leave out lone pairs on heteroatoms, but don’t forget about them. Use the formal charge to determine the number of lone pairs.

Orbitals and Bonding

• When the 1s orbital of one H atom overlaps with the 1s orbital of another H atom, a sigma (σ) bond that concentrates electron density between the two nuclei is formed. All single bonds are s bonds. This bond is cylindrically symmetrical because the electrons forming the bond are distributed symmetrically about an imaginary line connecting the two nuclei.
• To account for the bonding patterns observed in more complex molecules, we must take a closer look at how the 2s and 2p orbitals of atoms in the second row are utilized.
• In addition to its two core electrons, carbon has four valence electrons. In its ground state, carbon places two electrons in the 2s orbital and one each in 2p orbitals.
  • Note: The lowest energy arrangement of electrons for an atom is called its ground state.
• In this description, carbon should form only two bonds because it has only two unpaired valence electrons. However, the resulting species, CH2, is very unstable and cannot be isolated under typical laboratory conditions. Note that in CH2, carbon would not have an octet of electrons.
• To solve this dilemma, chemists have proposed that atoms like carbon do not use pure s and pure p orbitals in forming bonds. Instead, atoms use a set of new orbitals called hybrid orbitals.
• Hybridization is the combination of two or more atomic orbitals to form the same number of hybrid orbitals, each having the same shape and energy.
• The mixing of a spherical 2s orbital and three dumbbell shaped 2p orbitals together produces four hybrid orbitals, each having one large lobe and one small lobe.
• The four hybrid orbitals are oriented towards the corners of a tetrahedron, and form four equivalent bonds.
• Each bond in CH4 is formed by overlap of an sp3 hybrid orbital of carbon with a 1s orbital of hydrogen. These four bonds point to the corners of a tetrahedron.

Determining Hybridization

• Count the number of groups (atoms and nonbonded electron pairs) around the atom. The number of groups corresponds to the number of atomic orbitals that must be hybridized to form the hybrid orbitals.
  • Two groups: two sp hybrid orbitals
  • Three groups: three sp2 hybrid orbitals
  • Four groups: four sp3 hybrid orbitals
• Making a model of ethane illustrates one additional feature about its structure. Rotation occurs around the central C—C s bond.
• Unlike the C—C bond in ethane, rotation about the C—C double bond in ethylene is restricted. It can only occur if the π bond first breaks and then reforms, a process that requires considerable energy.
• Each carbon atom has two unhybridized 2p orbitals that are perpendicular to each other and to the sp hybrid orbitals.
• The side-by-side overlap of two 2p orbitals on one carbon with two 2p orbitals on the other carbon creates the second and third bonds of the triple bond. All triple bonds are composed of one sigma and two pi bonds.
• two bonds
  Csp-Csp
  C2p-C2p C2p-C2p

Bond Length and Bond Strength

• As the number of electrons between two nuclei increases, bonds become shorter and stronger. Triple bonds are shorter and stronger than double bonds, which are shorter and stronger than single bonds.
• The length and strength of C—H bonds vary depending on the hybridization of the carbon atom.

sp hybrid
one 2s orbital
two hybrid orbitals
= 50% s-character
sp² hybrid
one 2s orbital
three hybrid orbitals
= 33% s-character
sp³ hybrid
one 2s orbital
four hybrid orbitals
= 25% s-character
Increased percent s-character --- Increased bond strength --> Decreased bond length

Electronegativity

• Electronegativity is a measure of an atom’s attraction for electrons in a bond.
• Electronegativity values are used to indicate whether the electrons in a bond are equally shared or unequally shared between two atoms. When electrons are equally shared, the bond is nonpolar.
• A carbon—carbon bond is nonpolar. C—H bonds are considered to be nonpolar because the electronegativity difference between C and H is small.
• Whenever two different atoms having similar electronegativities are bonded together, the bond is nonpolar.
• Bonding between atoms of different electronegativity values results in unequal sharing of electrons.
• Example: In the C—O bond, the electrons are pulled away from C (2.5) toward O (3.4), the element of higher electronegativity. The bond is polar, or polar covalent. The bond is said to have dipole; that is, partial separation of charge. A C—O bond is a polar bond.

Depicting Polarity

• The δ+ means the indicated atom is electron deficient.
• The δ− means the indicated atom is electron rich.
• The direction of polarity in a bond is indicated by an arrow with the head of the arrow pointing towards the more electronegative element. The tail of the arrow is drawn at the less electronegative element.
• Use the following procedure to determine if a molecule has a net dipole:
  • Use electronegativity differences to identify all of the polar bonds and the directions of the bond dipoles.
  • Determine the geometry around individual atoms by counting groups, and decide if individual dipoles cancel or reinforce each other in space.
• A polar molecule has either one polar bond, or two or more bond dipoles that reinforce each other. An example is water:
  • It is a bent molecule
  • Two dipoles reinforce
  • It has a net dipole, making it a polar molecule
• A nonpolar molecule has either no polar bonds, or two or more bond dipoles that cancel. An example is carbon dioxide:
  • It is a linear molecule
  • Two dipoles are equal and opposite in direction
  • Two dipoles cancel
  • It is a nonpolar molecule with no net dipole