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A chemical bond is the force that holds atoms together in chemical compounds. There are two idealized types of bonding: (1) covalent bonding, in which electrons are shared between atoms in a molecule or polyatomic ion, and (2) ionic bonding, in which positively and negatively charged ions are held together by electrostatic forces.
- Band Structure
- Band Theory was developed with some help from the knowledge gained during the quantum revolution in science. In 1928, Felix Bloch had the idea to take the quantum theory and apply it to solids. In 1927, Walter Heitler and Fritz London discovered bands- very closely spaced orbitals with not much difference in energy.
- Bond Energies
- The bond energy is a measure of the amount of energy needed to break apart one mole of covalently bonded gases. Energy is released to generate bonds, which is why the enthalpy change for breaking bonds is positive. Energy is required to break bonds. Atoms are much happier when they are "married" and release energy because it is easier and more stable to be in a relationship (e.g., to generate octet electronic configurations).
- Bond Order and Lengths
- Bond order is the number of chemical bonds between a pair of atoms and indicatesthe stability of a bond. For example, in diatomic nitrogen,N≡N,the bond order is 3;in acetylene, H−C≡C−H, the carbon-carbon bond orderis also 3, and the C−H bond order is 1. Bond order and bond length indicate the type andstrength of covalent bonds between atoms. Bond order and length are inversely proportional to each other: when bond order is increased, bond length is decreased.
- Chemical Bonds
- Bond Lengths and Energies
- Chemical Bond
- Chemical Bonding - A Review
- Electron Density of \(\sigma\) and \(\pi\) Bonds
- Enthalpies of Reactions
- Ionic Compounds
- Lewis Dot Structures
- Molecular Orbitals of H2
- Valence Bond Theory and Hybrid Atomic Orbitals
- Contrasting MO and VB theory
- Both the MO and VB theories are used to help determine the structure of a molecule. Unlike the VB theory, which is largely based off of valence electrons, the MO theory describes structure more in depth by taking into consideration, for example, the overlap and energies of the bonding and antibonding electrons residing in a particular molecular orbital.
- Coordinate (Dative Covalent) Bonding
- A coordinate bond (also called a dative covalent bond) is a covalent bond (a shared pair of electrons) in which both electrons come from the same atom.
- Covalent Bonding
- This page explains what covalent bonding is. It starts with a simple picture of the single covalent bond.
- Covalent Bonds
- Covalent bonding occurs when pairs of electrons are shared by atoms. Atoms will covalently bond with other atoms in order to gain more stability, which is gained by forming a full electron shell. By sharing their outer most (valence) electrons, atoms can fill up their outer electron shell and gain stability. Nonmetals will readily form covalent bonds with other nonmetals in order to obtain stability.
- Covalent Bonds vs Ionic Bonds
- Covalent Bond Distance, Radius and van der Waals Radius
- Electrostatic Potential maps
- Electrostatic potential maps, also known as electrostatic potential energy maps, or molecular electrical potential surfaces, illustrate the charge distributions of molecules three dimensionally. These maps allow us to visualize variably charged regions of a molecule. Knowledge of the charge distributions can be used to determine how molecules interact with one another.
- Ionic Bonds
- Ionic bonding is the complete transfer of valence electron(s) between atoms and is a type of chemical bond that generates two oppositely charged ions. It is observed because metals with few electrons in its outer-most orbital. By losing those electrons, these metals can achieve noble-gas configuration and satisfy the octet rule. Similarly, nonmetals that have close to 8 electrons in its valence shell tend to readily accept electrons to achieve its noble gas configuration.
- Metallic Bonding
- A strong metallic bond will be the result of more delocalized electrons, which causes the effective nuclear charge on electrons on the cation to increase, in effect making the size of the cation smaller. Metallic bonds are strong and require a great deal of energy to break, and therefore metals have high melting and boiling points.A metallic bonding theory must explain how so much bonding can occur with such few electrons (since metals are located on the left side of the periodic table and do
- Non-Singular Covalent Bonds
- There are three types of covalent bonds: single, double, and triple. The name "Non-singular covalent bonds" speaks for itself. Non-singular covalent bonds are covalent bonds that need to share more then one electron pair, so they create double and triple bonds.
- Valence-Shell Electron-Pair Repulsion Models
- The valence-shell electron-pair repulsion (VSEPR) model is used to predict three-dimensional arrangements of atoms or bonds in molecules including bond lengths, bond angles and qualitative bond energies.
