ORGANIC CHEMISTRY I CHEM 2323

STRUCTURE and PROPERTIES

Organic Chemistry

Organic Chemistry - the chemistry of most compounds of carbon. Organic chemistry is a major division because of the possibility of multiple carbon arrangements in organic compounds.  Almost all known organic compounds   contain hydrogen.

The major source of the raw materials for industrial organic chemicals is the fossil fuels.

Organic chemistry is very dependent on structural theory.  By knowing the structure of a compound, we can predict physical properties (boiling and melting points, etc) and chemical behavior (reactions, solvents, etc)

Covalence numbers

Covalence numbers are the number of bonds an element will typically form.  Covalence numbers (the number of bonds) are used to draw the structures of organic compounds and determine molecular formulas.

C = 4  (CH₄)

N = 3  (NH₃)

H = 1  (HBr)

O and S = 2  (H₂S and H₂O

Halogens = 1  (HBr)

For example, carbon normally forms four bonds with other compounds (CH₄) while hydrogens only form one bond.

Atomic Orbitals

Atomic orbitals tell us the approximate probability of finding an electron at any particular place.  An orbital is a region in space where an electron is likely to be found.  Each electron spins about its axis in one of two directions.

An orbital can hold a maximum of two electrons.   When two electrons occupy an atomic orbital and to do so they must have opposite spins (Pauli exclusion principle).  This set of oppositely spinning electrons are called paired electrons.  The s orbital can hold only two electrons, but three p orbitals (p_x, p_y, p_z) with two electrons in each are needed for the six possible electrons in the p subshell. 

The s orbital is shaped like a sphere

The p orbitals are like barbells.

Covalent Bond

The vast majority of organic compounds are held together by covalent bonds.  Covalent bonds like their counterpart, ionic bonds, are driven by the need to form a noble gas like electron configuration (s²p⁶).   Covalent bonds are where two atoms share a pair of electrons.

The are three types of covalent bonds:

polar - where the electron pair is closer to one atom than the other.  This is due to the difference between the electronegativity of the two atoms.  Electronegativity is the tendency of the bonded atoms to attract electrons toward themselves.  The greater the difference, the higher the polarity.   The electronegativity increases on a periodic table from left to right (fluorine is greater than carbon) and from bottom to top (fluorine is greater than bromine).  NaF is an example of a polar compound.

nonpolar - the electron pair is shared equally between the two atoms (i.e. two atoms of similar electronegativity or two atoms of the same element - H₂)

coordinate - where one atom supplies both of the electrons in the pair.  NH₄⁺ is an example of a coordinate bond.  The nitrogen shares an unshared pair of electrons with a hydrogen that previously had no electrons.

In order for a covalent bond to form, two atoms must be located so that an atomic orbital of one overlaps the atomic orbital of the other.   One electron in each atomic orbital.

The two atomic orbitals merge to form a single bond orbital occupied by both electrons.   The two atomic orbitals merge to form one bond orbital occupied by both electrons.  These electrons must be paired.  The electrons belong to both nuclei. 

The two combined electrons now contain less energy (more stable) than the total of the original atomic orbitals.  The formation of the bond evolves energy.  This energy evolved is called the bond dissociation energy.

The greater the overlap of the two atomic orbitals - the stronger the bond.

Covalent bonds are strong because:

• in a single atom:  one p⁺ attracts one e^-

• but in a covalent molecule:  two p⁺ attract one e^-.

Bond length - distance between two nuclei.

The shape of the resulting orbital overlap is similar to the original atomic orbital shapes.  But the combination of different types of atomic orbitals, although similar in shape, will result in orbitals of characteristic length and strength.

Hybrid Orbitals

If two p orbitals overlapped we would expect the resulting angle between the two nuclei to be 90^o.

But angles of 180^o , 120^o and 109.5^o are observed in organic compounds instead of 90^o.    To account for these bond angles, Pauli suggested that the atomic orbitals may combine to form new orbitals which in turn interact to form bonds with the observed angles.  The combination of these atomic orbitals is called hybridization, and the resulting orbitals formed are called hybrid orbitals.

sp Hybrid Orbitals- linear (180^o)

sp hybrid orbitals are formed when an atom has a pair of 2s electrons but no 2p electrons.  Since the 2s electrons are already paired, no further bonding can take place.  To overcome this, one of the 2s electrons is promoted to the 2p level, thus producing two unpaired electrons available for bonding. 

Be	1s2	2s2	2px	2py	2pz
normal
hybrid

These hybrid orbitals are called sp since they are a combination of a s and a p orbital.  The atom now has two bonding orbitals.  In order for these two orbitals to get as far away from each other as possible along a flat plane, the bond angle for the overlap of these orbitals with others would be 180^o.

An example would be if Be combined with Cl.   The sp hybrid orbitals from Be would overlap the p orbitals of Cl.  Therefore, BeCl₂ would be a linear molecule with all three atoms aligned along a flat plane.

sp² Hybrid Orbitals - flat triangle (120^o)

sp² hybrid orbitals are formed when an atom has a pair of 2s electrons but only one 2p electron.   Since the 2s electrons are already paired, only one bonding electron is available. Atoms like to form as many bonds as possible.  Therefore, one of the 2s electrons is promoted to the 2p level, thus producing three unpaired electrons available for bonding. 

B	1s2	2s2	2px	2py	2pz
normal
hybrid

These hybrid orbitals are called sp² since they are a combination of a s and two p orbital.  The atom now has three bonding orbitals.  In order for these three orbitals to get as far away from each other as possible along a flat plane, the bond angle for the overlap of these orbitals with others would be 120^o.

An example would be if B combined with F.  The sp² hybrid orbitals from B would overlap the p orbitals of F.   Therefore, BF₃ would be a flat triangular molecule.

sp³ Hybrid Orbitals - tetrahedral (109.5^o)

sp³ hybrid orbitals are formed when an atom has a pair of 2s electrons but only two 2p electron.   Since the 2s electrons are already paired, only two bonding electrons are available.   Atoms like to form as many bonds as possible.  Therefore, one of the 2s electrons is promoted to the 2p level, thus producing four unpaired electrons available for bonding. 

C	1s2	2s2	2px	2py	2pz
normal
hybrid

sp³ orbitals are directed to the corners of a tetrahedron.  The angle between any two orbitals is the tetrahedral angle of 109.5^o.

Unshared pairs of electrons

Unshared pairs of electrons occupy more space than shared electrons.  In a tetrahedral, some of the sp³ orbitals may contain a pair of unshared electrons. 

Intramolecular forces

The actual structure of a molecule depends a great deal on the net results of all the repulsive and attractive forces within the molecule itself.

repulsive forces:

• e^-/e^-
• p⁺/p⁺
• two electrons spinning in same direction (paired)

attractive forces:

• e^-/p⁺
• two e^- spinning in opposite directions

Bond Dissociation Energy

Energy consumed or liberated when a bond is broken or formed.  Bond dissociation energies are characteristic of the type of bonds involved.  The breaking of a covalent bond to form homolytic fragments takes less energy than producing heterolytic fragments (the separation of charged particles takes more energy than separating neutral particles).

Polarity of Bonds

In covalent bonds, two nuclei share a pair of elections.  If the covalent electrons are not shared equally, it is called a polar bond.  One nuclei would have a partially negative charge (d^-), while the other nuclei would have a partial positive charge (d⁺).

Differences in electronegativity between the nuclei form polar bonds.  The great the differences in electronegativity, the more polar the bond.

Electronegativity: F > O > Cl , N > Br > C , H

Bond polarities play an important role in both physical and chemical properties.

Polarity of Molecules

Molecules are polar if the center of the partial negative charge does not coincide with the center of the partial positive charge.

Dipole - two equal and opposite charges separated in space.

Structure and Physical Properties

The type of bonds within the molecule can determine the structure and the resulting structure determines the physical properties (i.e. melting points, boiling points, etc).

Melting Point

Solids have a highly ordered arrangement of particles.  Melting is when the highly ordered structure becomes more random.   When enough energy is supplied for the thermal energy of the particles to overcome the intracrystalline forces, melting occurs.  The temperature at which melting occurs is called the melting point (MP).

Ionic compounds -

• structures units are atoms
• interionic forces are strong
• high melting points

Non-ionic compounds -

• structured units are molecules
• intermolecular forces are weak
• low melting points.

Intermolecular Forces

Forces which hold neutral molecules to each other appear to be electrostatic in nature.  There is an attraction between the positive charges and negative charges.  There are two types of intermolecular forces:

• Dipole - dipole interactions (hydrogen bonding) is the attraction of the partial positive charge of one molecule to the partial negative charge of another molecule.  Hydrogen bonding is a special case (common) where the hydrogen atom (partial positive) serves as a bridge between two electronegative atoms.  One atom is covalently bonded to the hydrogen, the other is electrostatically attracted.  For hydrogen bonding to be sufficient, both electronegative atoms must be a any combination or pairing of F, O , or N.

• Van der Waals forces - even though the atoms have no net dipole, the electrons move around a little bit forming "momentary" dipoles. Therefore, there can be a weak attraction.

Solubility

When a solid or liquid dissolves, the structural units become separated and the spaces in-between these structural units become occupied by solvent molecules.

In dissolving, energy must be supplied to overcome the old interionic or molecular forces.

The energy to break these bonds is supplied by the formation of bonds between the solute particles and solvent particles.

The old attractive forces between the structural units are replaced by new attractive forces between the solvent and the structural units..

Boiling Point

In liquids, the particles are less regular in structure and have freedom of movement, but each particle is still attracted to a number of other particles.  Boiling occurs when enough thermal energy is supplied to overcome these attractive forces.  Molecules break away from each other to form gases.

Ionic bonds are strong, therefore have high boiling point.

Covalent bonds are weak, therefore have lower boiling point.

As the molecular weight increases, the boiling point and melting point increases.

At temperatures greater than 350^oC, covalent compounds start to break down internally (decompose).

Isomerism

Isomers are compounds with the same molecular formula but different structures.  The different molecular structures cause differences in chemical and physical properties.