4.1 Types of Chemical Bonds

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• Electrostatic forces that hold atoms together in compounds
• Involves the interaction of the valance electrons

<<Ionic<< <<Chemical Bonds<<

• electrons and transferred between oppositely charged ions
• form ionic compounds with ionic bonding
• ions form to become isoelectronic with a noble gas, same electron configuration   * Isoelectronic: same electron configuration as another

<<Covalent Chemical Bonds<<

• electrons are shared
• forms molecular compounds
• Nonpolar, polar or pure
• Molecular element: pure substance made up of 2 atoms like Oxygen
• As two atoms move close together, the electron cloud of one attracts the nucleus of the other
• At the same time the nuclei repel each other as do the electron cloud
• The atoms stay a distance from one another that has the lowest overall energy of the system like the hydrogen molecule.
• Inter: how the molecules are attracting with one another like London Dispersion, Hydrogen Bonding & Dipole-Dipole
• Intra: what is holding the molecule togehter like Covalent or Ionic

<<Lewis Theory of Bonding<<

• Atoms & ions are stable if they have a full valence shell of electrons
• Electrons are most stable when they are paired
• Atoms form chemical bonds to achieve a full valence shell of electrons
• A full valence shell of electrons may be achieved by an exchange of electrons between metal and non-metal atoms
• The sharing of electrons results in a covalent bond

<<Duet / Octet Rule<<

• Hydrogen is stable with 2 electrons
• Most atoms are stable with 8 electrons

<<Lewis Structures<<

1. Draw the central atom (highest bonding capacity)
2. Arrange the symbols of the atoms for the rest of the elements around equal distance apart
3. Add up the number of valence electrons of each atom. Add to this any negative charge or subtract any positive charge.
4. Place a pair of bonded electrons between central atom and each of the others (single bond).
5. Place lone electron pairs on outer atoms first (follow duet and octet rule)
6. Dump rest of electrons on central atom in pairs
7. Move electrons around to form double or triple bonds until all atoms follow octet/duet rule.

• Resonance Structures: models that give the relative position of atoms in a Lewis Structure, but show different places for their bonding and lone pairs

Exceptions

• Under filled octets: molecules who’s central atoms are surrounded by fewer than 8 electrons   * Ex: Boron Trifluoride

• Overfilled octets: molecules whose central atoms are surrounded by more than 8 electrons   * Sulfur has its valence electrons in the third energy level   * There is space for 12 electrons in the valence shells   * Sulfur, phosphorus, chlorine

Determining Formal Charge

• Formal charge of a molecule to determine its most stable structure

• Determine for each atom, then add all atoms together to determine charge of molecule
• The arrangement with which charges are the most stable

• The arrangement where P has an overfilled octet contains fewer charges than the arrangement where P has a normal octet
• Certain elements are able to do this because they have an empty 3d orbital that electrons can fill

Coordinate Covalent Bonding

• a covalent bond in which the both of the bonding electrons are from one atom

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