Pharmaceuticals

Covalent Bonding:

• Electrons are shared rather than transferred

• Can be be simple with a molecule structure (H₂O) or can be a complex covalent structure (Graphite)

Octet Shell Exceptions

• H, He, Li, Be do not obey this rule as they exist with 2 electrons in their outermost energy level

• Boron can exist both as 2 and 8 electrons but most stable is 6 electrons

Bond length reason

• The bonding electrons draw the atoms closer together but as the atoms get very close they experience repulsive forces from the other non-bonding and from the two nuclei themsleves

• This means that there is a minimum distance for the two atoms which called the bond length

• Generally as bond strength increases bond length decreases

Coordinate Bonding

. In a coordinate covalent bond one of the atoms contributes more electrons than the other. A common example of this is the bond in carbon monoxide.

Three Postulates of VSEPR

• Valence electrons are invovled in bonding

• Lone electrons pairs and shared electrons pairs on the central atom

• Lone pairs repel more than shared pairs, decreasing bonding angle by 2.5^o

2 Electron Domains,

Linear shape, 180^o

Electron domains 3

Trigonal planar, 120^o

Electron domains 3, lone pair 1

Trigonal planar, Bent shape, 117.5^o

Electron domains 4

Tetrahedral, 109.5

Electron domains 4, lone pair 1

electron domain geometry: tetrahedral molecular geometry: Trigonal pyramidal, 107

Electron Domains 4, lone pairs 2

Tetrahedral, Bent shape104.5^o

Intermolecular forces strength (increasing order)

dispersion forces, dipole dipole forces, hydrogen bonds, covalent bonds, ion dipole forces

Ionic attraction compounds

The attraction between the ions is an electron static attraction and seen as an intramolecular forces

Polar covalent molecules intermolecular force

Intermolecular forces will be mainly dipole dipole and may be hydrogen if the hydrogen in the molecule is bonded to a very electronegative oxygen, nitrogen or fluorine

Force in non polar Covalent Molecules

Intermolecular forces only London Dispersion Forces

London dispersion forces result from a temporarily induced dipole

Uneven movement of electrons can result in a bunch of electrons located on one side of the molecule which gives that side of the molecule a slightly negative charge. A molecule near it will have a positive charge induced on it.

Increase in strength with higher molar mass as increased number of electrons and protons, weaker than hydrogen bonding and covalent bonding.

Diamond

Giant 3D covalent structure each carbon is bonded to four other carbons in a tetrahedral shape

sp3

Graphite

Giant 2D covalent structure

Has delocalized electrons so can transfer current

sp2

Each graphite has three bonds

Fullerenes

C60 composed of 5 tor 6 cabon rings

Exactly the same as graphite except a single sheet

sp2

Sigma bond:

Formed by the direct overlap of orbitals

Backbone of every molecule, every covalent bond has a sigma bond. Stronger than pi bonds due to closer to nuclei

Formal Charge

Hypothetical charge molecule would have based on the number of electrons are bonded to it

Key is to make the charges as small as possible

Expanded octets

Atoms in the 3rd period are able to use d-orbitals , which allows them to form more bonds than the octet rule predicts

The same rule apply, except that the central atom will have more bonds than you would thought it was allowed to have

Pi bonds

sideways overlap unhybridized orbitals

These form the double or triple bonds

Double bond 1σ and 1π

Triple bond 1σ and 2π bonds

Less strong because farther away from the nuclei

Hybridisation

The mixing of atomic orbitals to produce degenerate hybrid orbitals to be used for bonding.

sp3 hybridazation

tetrahedral, 4 electron domain, 1s+3p

sp2 hybridization

Trigonal planar, 3 electron domains, 1s+2p

sp

linear, 2 electron domains, 1s+1p

5 electron domains

electron domain geometry: trigonal bipyramidal

Molecular domain geometry: trigonal bipyramidal 90,120

6 electron domains

electron domain geometry: octahedral

molecular geometry: octahedral 90

5 electron domains and one lone pair

electron domain geometry: trigonal bipyramidal

Molecular domain geometry: see-saw ~90,~120

5 electron domains and two lone pairs

electron domain geometry: trigonal bipyramidal

Molecular domain geometry: T-Shape 90 and 180

5 electron domains and three lone pairs

electron domain geometry: trigonal bipyramidal

Molecular domain geometry: linear 180

Electron domain geometry

Position of all electron groups

Molecular geometry

Only the position of connected atoms, same as electron domain geometry if no lone pairs

Electron domain 6 1 lone pair

electron domain geometry: octahedral

molecular geometry: square pyramidal 87.5

Electron domain 6 2 lone pairs

electron domain geometry: octahedral

molecular geometry: square planar 90

Electron domain 6 3 lone pairs

electron domain geometry: octahedral

molecular geometry: T shaped 90 and 180

Electron domain 6 4 lone pairs

electron domain geometry: octahedral

molecular geometry: linear 180

How bonds from s orbitals form

s orbital: spherical

Forms bonds by head-on overlap

Can overlap with s or p orbitals

Forms a σ (sigma) bond

What p orbitals form

p orbital: two lobes

Head-on overlap → σ bond

Sideways overlap → π bond

π bonds require parallel p orbitals

resonance

Resonance when a molecule has two or more valid Lewis structures that differ only in the arrangement of electrons, while the positions of the atoms stay the same. For example Ozone there is not one single bond and one double bond, both have same length, electrons switch between them

Benzene

Alternating three single bonds and three double bonds

C6H6

1. Does not undergo electrophilic addition reactions

2. Bond strength is the same

3. The bond length is the same

Due to extra bond strength and bond length extra stability

Structural isomers

Same molecular formula but different structural formula

cis/trans and E/Z isomers

position about a carbon double bond. Has to be broken in order to convert between different isomers, retriction in rotation due to pi bond.

Cis isomer

similar parts are on the same side

trans isomers

similar parts are on different sides

Optical isomers (enantionmers)

• Optical isomers have chiral centres - element attached to four different substituent groups

• Optical isomers ome in pairs they are mirror images of each other

• Optical isomers rotate plane polarized light otherwise have the same chemical and physical properties

• Both isomers rotate light the same amount but in different directions

Diasteromers

contain two or more chiral centres. They are non imposable but do not form mirror images. Diastereomers have different physical and chemical properties

racemic mixture

Equal amount of left and right handed enantiomers

Stereo isomers

Same chemical formula and atom bonding formula but differ how they are arranged in 3d space

Silicon dioxide

Giant covalent network where silicon is sp3 hybridized and oxygen is sp3 hybridized.

The H NMR spectrum of an organic compound provides information concerning

• the # of different types of hydrogens present in the molecule

• the relative #’s of the different types of hydrogens

• the electronic environment of the different types of hydrogens

• the number of hydrogen “neighbor”’ a hydrogen has

IR numbers

reciprocals of wavelength (1/λ). They are also referred to as wavenumbers and measured in cm-1

Infrared spectra

Lighter atoms will vibrate at higher frequencies and heavier atoms at lower frequencies.

In order for a covalent bond to absorb infrared radiation there must be change in the dipole moment. Diatomic molecules containing only one element eg H₂ do not absorb infrared radiation

IR spectrums can be used to identify functional groups.

Chiral center

an atom in a molecule bonded to four different groups, creating a structure that cannot be placed on top of its mirror image, enantionmer

Allotrope

Different forms of the same element / different (bonding) arrangements of the atoms of an element in the same physical state.

Nucleophile

negative thus nucleus/proton seeking

Electronphile

positive thus electron seeking

singlet

Multiplicity (number of peaks H NMR)

Tells how many equivalent hydrogen atoms are coupling with hydrogen observed. Happens due to spin to spin coupling between neighbouring hydrogen nuclei. Follows n+1 rule

Integration

Sum of hydrogens producing each signal. Add them if chemically equivalent, area under a signal.

Curly arrows

A curly arrow is used to show the direction of electrons within bonds, drawn from the site of electron availability, such as a pair of lone electrons, to the site of electron deficiency, such as an atom with a partial positive charge

Free radicals

Unstable and reactive atoms with at least one lone unpaired electron in their valence shell

Electrophilic addition

An addition reaction where a π bond in an unsaturated molecule (such as an alkene) is broken by an attack from an electrophile to form two new σ bonds. Ethene and HCL

Addition polymerisation

A reaction where unsaturated monomers join together without forming by-products to form a polymer.

Nucleophilic substitution

A reaction where an electron-rich nucleophile replaces a leaving group in a molecule.

Free-radical substitution

A substitution reaction involving free radicals, proceeding through initiation, propagation, and termination steps.

Heterolytic fission

the unequal breaking of a covalent bond where one bonded atom keeps both shared electrons, forming a cation and an anion

Lewis Acid

a substance that accepts a pair of electrons to form a coordinate covalent bond. Any substance that has an incomplete octet will accept an electron pair.

Lewis Base

a substance that donates a pair of electrons to form a coordinate covalent bond. Any substance that has a lone pair.