1: Structure and Bonding
- Page ID
- 418057
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Much of the material in chapter one reviews topics covered CHM 115, while applying those concepts to organic molecules.
- the differences between organic and inorganic chemistry
- ionic and covalent bonding
- polar and nonpolar bonds
- Lewis structures
- hybridization
- molecular geometry and dipole moment
- the structure and geometry of the compounds methane, ethane, ethylene and acetylene.
- 1.1: Describing Chemical Bonds - Valence Bond Theory
- Covalent bonds form as valence electrons are shared between two atoms. Lewis Structures and structural formulas are common ways of showing the covalent bonding in organic molecules. Formal charge describes the changes in the number of valence electrons as an atom becomes bonded into a molecule. If the atom has a net loss of valence electrons it will have a positive formal charge. If the atom has a net gain of valence electrons it will have a negative formal charge.
- 1.2: Describing Chemical Bonds- Molecular Orbital Theory
- Molecular Orbital theory (MO) is a more advanced bonding model than Valence Bond Theory, in which two atomic orbitals overlap to form two molecular orbitals – a bonding MO and an anti-bonding MO.
- 1.3: Polar Covalent Bonds - Electronegativity
- Because the tendency of an element to gain or lose electrons is so important in determining its chemistry, various methods have been developed to quantitatively describe this tendency. The most important method uses a measurement called electronegativity, defined as the relative ability of an atom to attract electrons to itself in a chemical compound.
- 1.4: Development of Chemical Bonding Theory
- Lewis Dot Symbols are a way of indicating the number of valence electrons in an atom. They are useful for predicting the number and types of covalent bonds within organic molecules. The molecular shape of molecules is predicted by Valence Shell Electron Pair Repulsion (VSEPR) theory. The shapes of common organic molecules are based on tetrahedral, trigonal planar or linear arrangements of electron groups.
- 1.5: Formal Charges
- A formal charge is the charge assigned to an atom in a molecule, assuming that electrons in all chemical bonds are shared equally between atoms, regardless of relative electronegativity.
- 1.6: Drawing Chemical Structures
- Kekulé Formulas or structural formulas display the atoms of the molecule in the order they are bonded. Condensed structural formulas show the order of atoms like a structural formula but are written in a single line to save space. Skeleton formulas or Shorthand formulas or line-angle formulas are used to write carbon and hydrogen atoms more efficiently by replacing the letters with lines. Isomers have the same molecular formula, but different structural formulas
- 1.7: sp³ Hybrid Orbitals and the Structure of Methane
- The four identical C-H single bonds in methane form as the result of sigma bond overlap between the sp3 hybrid orbitals of carbon and the s orbital of each hydrogen.
- 1.8: sp³ Hybrid Orbitals and the Structure of Ethane
- The C-C bond in ethane forms as the result of sigma bond overlap between a sp³ hybrid orbital on each carbon. and the s orbital of each hydrogen. The six identical C-H single bonds in form as the result of sigma bond overlap between the sp³ hybrid orbitals of carbon and the s orbital of each hydrogen.
- 1.9: sp² Hybrid Orbitals and the Structure of Ethylene
- The C=C bond in ethylene forms as the result of both a sigma bond overlap between a sp2 hybrid orbital on each carbon and a pi bond overlap of a p orbital on each carbon
- 1.10: sp Hybrid Orbitals and the Structure of Acetylene
- The carbon-carbon triple bond in acetylene forms as the result of one sigma bond overlap between a sp hybrid orbital on each carbon and two pi bond overlaps of p orbitals on each carbon.
- 1.11: Hybridization of Nitrogen, Oxygen, Phosphorus and Sulfur
- The atomic orbitals of nitrogen, oxygen, phosphorus and sulfur can hybridize in the same way as those of carbon.
- 1.12: Polar Covalent Bonds - Dipole Moments
- Mathematically, dipole moments are vectors; they possess both a magnitude and a direction. The dipole moment of a molecule is therefore the vector sum of the dipole moments of the individual bonds in the molecule. If the individual bond dipole moments cancel one another, there is no net dipole moment.