Introduction
“Organic” – until mid 1800’s referred to compounds from living sources (mineral sources were“inorganic”)
“Wöhler” in 1828 showed that urea, an organic compound, could be made from a minerals.Today, organic compounds are those based on carbon structures and organic chemistry studies their structures and reactions
1.Includes biological molecules, drugs, solvents, dyes
2.Does not include metal salts and materials (inorganic)
3.Does not include materials of large repeating molecules without sequences (polymers)
Task
1. Introducing the Atomic Structure:
1)Structure of an atom
2)Electron Configurations
3)The Nature of the Chemical Bond
2. Giving some examples about the Valences of some matters:
1)Valences of Carbon
2)Valences of Oxygen
3)Valences of Nitrogen
3. Explaining the Non-bonding electrons and Valence Bond Theory
4. Conclusion
Process
1. Introducing the Atomic Structure:
1)Structure of an atom
1. Positively charged nucleus (very dense, protons and neutrons) and small (10-15 m).
2. Negatively charged electrons are in a cloud (10-10 m) around nucleus
3. Diameter is about 2 × 10-10 m (200 picometers (pm)
2)Electron Configurations
Ground-state electron configuration of an atom lists orbitals occupied by its electrons. Rules:
1. Lowest-energy orbitals fill first: 1s → 2s → 2p → 3s → 3p → 4s → 3d
2. Electron spin can have only two orientations, up ↑ and down ↓. Only two electrons can occupy an orbital, and they must be of opposite spin (Pauli exclusion principle) to have unique wave equations.
3. If two or more empty orbitals of equal energy are available, electrons occupy each with spins parallel until all orbitals have one electron (Hund's rule).
3)The Nature of the Chemical Bond
1. Atoms form bonds because the compound that results is more stable than the separate atoms
2. Ionic bonds in salts form as a result of electron transfers
3. Organic compounds have covalent bonds from sharing electrons (G. N. Lewis, 1916)
4. Lewis structures shown valence electrons of an atom as dots
5. Hydrogen has one dot, representing its 1s electron Carbon has four dots (2s2 2p2)
6. Stable molecule results at completed shell, octet (eight dots) for main-group atoms (two for hydrogen)
2. Giving some examples about the Valences of some matters:
1)Valences of Carbon Carbon has four valence electrons (2s2 2p2), forming four bonds (CH4)
2)Valences of Oxygen Oxygen has six valence electrons (2s2 2p4) but forms two bonds (H2O)
3)Valences of Nitrogen Nitrogen has five valence electrons (2s2 2p3) but forms only three bonds (NH3)
3. Explaining the Non-bonding electrons and Valence Bond Theory
1). Non-bonding electrons
a. Valence electrons not used in bonding are called nonbonding electrons, or lone-pair electrons.
b. Nitrogen atom in ammonia (NH3) Shares six valence electrons in three covalent bonds and remaining two valence electrons are nonbonding lone pair.
2).Valence Bond Theory
a. Covalent bond forms when two atoms approach each other closely so that a singly occupied orbital on one atom overlaps a singly occupied orbital on the other atom
b. Electrons are paired in the overlapping orbitals and are attracted to nuclei of both atoms
c. H–H bond results from the overlap of two singly occupied hydrogen 1s orbitals
d. H-H bond is cylindrically symmetrical, sigma (σ) bond
Evaluation
There are many tests about organic chemistry which are pretty good for stundents to practice . The author recommend students to do these tests in order to check that do they really understand the concepts.
http://www.chemhelper.com/practicetests.html
Conclusion
Organic chemistry – chemistry of carbon compounds
Atom:
1. Positively charged nucleus surrounded by negatively charged electrons
2. Electronic structure of an atom described by wave equation
3. Electrons occupy orbitals around the nucleus.
Covalent bonds - electron pair is shared between atoms
Valence bond theory - electron sharing occurs by overlap of two atomic orbitals
Sigma (σ) bonds - Circular cross-section and are formed by head-on interaction