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1: Structure and Bonding

Last updated

Sep 20, 2022
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- About the Authors
- 1.0: Introduction to Organic Chemistry

Steven Farmer & Dietmar Kennepohl
LibreTexts

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Chapter Objectives

This chapter provides a review of material covered in a standard freshman general-chemistry course through a discussion of the following topics:

the differences between organic and inorganic chemistry.
the shapes and significance of atomic orbitals.
electron configurations.
ionic and covalent bonding.
molecular orbital theory.
hybridization.
the structure and geometry of the compounds methane, ethane, ethylene and acetylene.

1.0: Introduction to Organic Chemistry
Organic compounds contain carbon atoms bonded hydrogen and other carbon atoms. Organic chemistry studies the properties and reactions of organic compounds.
1.1: Atomic Structure - The Nucleus
Atoms are comprised of protons, neutrons and electrons. Protons and neutrons are found in the nucleus of the atom, while electrons are found in the electron cloud around the nucleus. The relative electrical charge of a proton is +1, a neutron has no charge, and an electron’s relative charge is -1. The number of protons in an atom’s nucleus is called the atomic number, Z. The mass number, A, is the sum of the number of protons and the number of neutrons in a nucleus.
1.2: Atomic Structure - Orbitals
This section explains atomic orbitals, emphasizing their quantum mechanical nature compared to Bohr's orbits. It covers the order and energy levels of orbitals from 1s to 3d and details s and p orbital shapes. It describes nodal planes, electron probability, and radial probability density. Differences between models by Bohr and Schr??dinger are highlighted, noting that orbitals are probability distributions rather than fixed paths.
1.3: Atomic Structure - Electron Configurations
The text aims to teach how to write ground-state electron configurations for elements up to atomic number 36, focusing on the arrangement of electrons in atomic orbitals. It explains key concepts such as electron configurations, Hund's rule, the Pauli exclusion principle, and the Aufbau principle. The periodic table is vital for determining these configurations, and rules for assigning electron orbitals are highlighted.
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: 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.
1.6: sp³ Hybrid Orbitals and the Structure of Methane
The text explains the structure of methane (CH4) using the concept of sp3 hybridization of the central carbon atom. Methane exhibits a tetrahedral shape with an H-C-H bond angle of 109.5??. Valence bond theory and Linus Pauling's hybridization model are highlighted to explain how carbon's 2s and 2p orbitals combine to form four equivalent sp3 orbitals. These hybrid orbitals allow the formation of four identical C-H sigma (??) bonds, accounting for methane's observed structure and bond properties.
1.7: 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.
1.8: sp² Hybrid Orbitals and the Structure of Ethylene
This section explains the formation of carbon-carbon double bonds through sp2 hybridization, resulting in sigma (??) and pi (??) bonds. In ethylene (CH2CH2), each carbon atom forms three sp2 hybrid orbitals for ?? bonds and has one unhybridized pz orbital that overlaps to form a ?? bond. This combination results in a trigonal planar geometry and makes ethylene's structure rigid due to the ?? bond restricting rotation. Overall, ethylene comprises five sigma and one pi bond.
1.9: sp Hybrid Orbitals and the Structure of Acetylene
This section covers the application of sp hybridization to explain carbon-carbon triple bonds, addressing the formation of sigma and pi bonds in such arrangements. It highlights the key characteristics of carbon bonds, including typical bond lengths and bond strengths associated with single, double, and triple bonds. A comparison is drawn between ethane, ethylene, and acetylene, showing that bond strength increases and bond length decreases with higher bond order.
1.10: Hybridization of Nitrogen, Oxygen, Phosphorus and Sulfur
This section explores the concept of hybridization for atoms like nitrogen, oxygen, phosphorus, and sulfur, explaining how these atoms form structures in simple compounds. The hybridization process involves mixing atomic orbitals to form new hybrid orbitals suitable for the pairing of electrons. For instance, nitrogen in ammonia (NH3) and methylamine, as well as oxygen in water (H2O) and methanol, exhibit sp3 hybridization.
1.11: 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.
1.12: 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

Chapter Objectives

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