Week 1.1 – fundamentals
Organic compounds contain the element carbon.
Atomic structure:
1. Nucleus
- Positively charged
- Protons and neutrons
2. Electron cloud
- Negatively charged
- Electrons
Orbital: describes volume of space around the nucleus that the electron is most likely to
occupy.
Orbitals contain 0,1 or 2 electrons,
and are grouped in electron shells of
increasing size/ energy,
There are three equal, mutually
perpendicular p orbitals: px, py, and pz
Lobes of p orbitals are separated by
region of zero electron density: a
node.
,Electron shells:
1st: S orbital of 2 electrons
2nd: S orbital and 3 P orbitals of 8 electrons
3rd: S orbital and 3 P orbitals and 5 d orbitals of 18 electrons
Electron configuration:
Most stable, lowest-energy electron configuration of an atom.
3 rules:
1. Lowest-energy orbitals fill first: 1s → 2s → 2p → 3s → 3p → 4s → 3d
2. Electron spin can have only 2 orientations: up and down
- A maximum of two electrons can occupy an orbital, and they must be of opposite
spin
3. If two or more empty orbitals of equal energy are available, electrons occupy each
orbital with parallel spins until all orbitals have one electron
Chemical bonding:
Atoms bond because the compound that results is more stable and lower in energy than the
separate atoms.
- Energy is released from the chemical system when a bond is formed.
- Energy is consumed by the system when a bond breaks.
1. Outer electron shell is the valence shell: in periodic table the row/ group number.
- Eight electrons in valence shell (an electron octet) impart special stability to noble-
gas elements.
- Other elements have a tendency to take on electron configuration of the nearest
noble gas.
2. Ionic compounds:
- Ion is formed when an electron is gained or lost from a neutral atom achieve an
octet configuration by gaining or losing electrons.
- Ions are held together by electrostatic attraction forming an ionic bond.
3. Covalent compounds: Molecule = Neutral collection of atoms held together by covalent
bonds.
- Bond formed by sharing electrons between atoms.
- Carbon achieves an octet configuration by sharing electrons.
Structures:
1. Lewis structure
- Valance shell electrons of an atom are represented as dots
2. Kekulé structure
- Two electron covalent bond is represented by a line
3. Lone- pair electrons
, - Valance shell electron pairs not used for bonding.
! Octet rule: Maximum of 8 electrons (4 electron-pairs) around the atoms C, N, O, F!
Number of covalent bonds depends on how many additional valence electrons needed to
reach noble-gas configuration.
Formal charges: if the number of bonds is different than the usual amount
- Positive (+) when the atom has one electron less than usual
- Negative (-) when the atom has one electron more than usual
- Neutral when the atom has correct amount of electrons so # of valance electrons
Chemical drawing structures:
Condensed structures: Carbon-hydrogen
and carbon-carbon bonds are not explicitly
shown.
Skeletal structure (line-bond structure):
Carbon atoms are assumed at each
intersection of two lines (bonds) and at the
or
Hybridization: mixing of orbitals
Valence bond theory:
- describes a covalent bond as resulting from the overlap of two atomic orbitals.
- Electrons are paired in overlapping orbitals and are attracted to nuclei of both
atoms, thus bonding the two atoms together.
VSERP model: valence shell electron-pair
repulsion
- All electron pairs (bonds and lone
pairs) surrounding an atom in the
valence shell will repel each other.
Geometry: with maximum distanced
between the groups of electrons.
C-C double bond is shorter and stronger than
C-C single bond
methane:
- Carbon has 4 valence electrons (2s22p2) to form 4 bonds.
- All C-H bonds in methane are identical and are spatially oriented toward the corners
of a regular tetrahedron (VSEPR).
- Single bond = σ bond
, Ethyne = acetylene (C2H2):
- Linear
- C-C triple bond, 6 shared electrons
- 180
- triple bond = σ bond + 2 ϖ bond
Intramolecular forces:
Covalent interactions:
polar bonds:
- A covalent bond in which the electron distribution between atoms is unsymmetrical,
due to difference in electronegativity (EN).
- EN: The ability of an atom to attract shared electrons in a covalent bond.
Polar substances ( zero) are more
soluble in polar solvents like water
(hydrophilic) or methanol.
apolar is hydrophobic, absence of polar
bonds or polar bonds that are opposite
they cancel each other out.
Intermolecular forces:
Noncovalent interactions:
1. Hydrogen bonds: A weak attraction between a hydrogen atom bonded to an
electronegative O or N and an electron lone pair on another O or N atom.
2. Van der Waals forces: molecule bond
3. Dipole- dipole forces : Occur between polar molecules as a result of electrostatic
interactions among dipoles.
Organic compounds contain the element carbon.
Atomic structure:
1. Nucleus
- Positively charged
- Protons and neutrons
2. Electron cloud
- Negatively charged
- Electrons
Orbital: describes volume of space around the nucleus that the electron is most likely to
occupy.
Orbitals contain 0,1 or 2 electrons,
and are grouped in electron shells of
increasing size/ energy,
There are three equal, mutually
perpendicular p orbitals: px, py, and pz
Lobes of p orbitals are separated by
region of zero electron density: a
node.
,Electron shells:
1st: S orbital of 2 electrons
2nd: S orbital and 3 P orbitals of 8 electrons
3rd: S orbital and 3 P orbitals and 5 d orbitals of 18 electrons
Electron configuration:
Most stable, lowest-energy electron configuration of an atom.
3 rules:
1. Lowest-energy orbitals fill first: 1s → 2s → 2p → 3s → 3p → 4s → 3d
2. Electron spin can have only 2 orientations: up and down
- A maximum of two electrons can occupy an orbital, and they must be of opposite
spin
3. If two or more empty orbitals of equal energy are available, electrons occupy each
orbital with parallel spins until all orbitals have one electron
Chemical bonding:
Atoms bond because the compound that results is more stable and lower in energy than the
separate atoms.
- Energy is released from the chemical system when a bond is formed.
- Energy is consumed by the system when a bond breaks.
1. Outer electron shell is the valence shell: in periodic table the row/ group number.
- Eight electrons in valence shell (an electron octet) impart special stability to noble-
gas elements.
- Other elements have a tendency to take on electron configuration of the nearest
noble gas.
2. Ionic compounds:
- Ion is formed when an electron is gained or lost from a neutral atom achieve an
octet configuration by gaining or losing electrons.
- Ions are held together by electrostatic attraction forming an ionic bond.
3. Covalent compounds: Molecule = Neutral collection of atoms held together by covalent
bonds.
- Bond formed by sharing electrons between atoms.
- Carbon achieves an octet configuration by sharing electrons.
Structures:
1. Lewis structure
- Valance shell electrons of an atom are represented as dots
2. Kekulé structure
- Two electron covalent bond is represented by a line
3. Lone- pair electrons
, - Valance shell electron pairs not used for bonding.
! Octet rule: Maximum of 8 electrons (4 electron-pairs) around the atoms C, N, O, F!
Number of covalent bonds depends on how many additional valence electrons needed to
reach noble-gas configuration.
Formal charges: if the number of bonds is different than the usual amount
- Positive (+) when the atom has one electron less than usual
- Negative (-) when the atom has one electron more than usual
- Neutral when the atom has correct amount of electrons so # of valance electrons
Chemical drawing structures:
Condensed structures: Carbon-hydrogen
and carbon-carbon bonds are not explicitly
shown.
Skeletal structure (line-bond structure):
Carbon atoms are assumed at each
intersection of two lines (bonds) and at the
or
Hybridization: mixing of orbitals
Valence bond theory:
- describes a covalent bond as resulting from the overlap of two atomic orbitals.
- Electrons are paired in overlapping orbitals and are attracted to nuclei of both
atoms, thus bonding the two atoms together.
VSERP model: valence shell electron-pair
repulsion
- All electron pairs (bonds and lone
pairs) surrounding an atom in the
valence shell will repel each other.
Geometry: with maximum distanced
between the groups of electrons.
C-C double bond is shorter and stronger than
C-C single bond
methane:
- Carbon has 4 valence electrons (2s22p2) to form 4 bonds.
- All C-H bonds in methane are identical and are spatially oriented toward the corners
of a regular tetrahedron (VSEPR).
- Single bond = σ bond
, Ethyne = acetylene (C2H2):
- Linear
- C-C triple bond, 6 shared electrons
- 180
- triple bond = σ bond + 2 ϖ bond
Intramolecular forces:
Covalent interactions:
polar bonds:
- A covalent bond in which the electron distribution between atoms is unsymmetrical,
due to difference in electronegativity (EN).
- EN: The ability of an atom to attract shared electrons in a covalent bond.
Polar substances ( zero) are more
soluble in polar solvents like water
(hydrophilic) or methanol.
apolar is hydrophobic, absence of polar
bonds or polar bonds that are opposite
they cancel each other out.
Intermolecular forces:
Noncovalent interactions:
1. Hydrogen bonds: A weak attraction between a hydrogen atom bonded to an
electronegative O or N and an electron lone pair on another O or N atom.
2. Van der Waals forces: molecule bond
3. Dipole- dipole forces : Occur between polar molecules as a result of electrostatic
interactions among dipoles.
