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Bonding - Coggle Diagram
Bonding
Ionic and covalent Bonding (L1)
Covalent Bonding
Properties
Bonds/Lewis dot Diagram: Covalent Bonds share electrons. When drawing the Lewis Dot diagram, Electrons will be shared to create LP between atoms.
Ionic Bonding
Properties
Bonds/Lewis Dot Diagram: Ionic bonds give electrons. When drawing lewis dot diagrams, one atom is trying to empty their valence shell, while the other is trying to fill it.
Terminology
Octet Rule: When atoms have the tendency to fill the outer valence shell with 8 electrons. Keep in mind that in some rare instances, there are exceptions to this rule and an atom may exceed 8 electrons.
BE/LP: Bonding Electrons and Lone Pairs
Crystal Lattice:
Network Solids:
Metallic Bonding:
Diagrams:
Lewis Dot:
Structural:
Reactivity and Electronegativity (L4):
Electronegativity: The ability of an atom to attract bonded electrons to itself within a covalent bond. (How strong is an atom at a game of tug of war)
Dipole: Is a positive pole and negative pole created with electronegativity. It is an uneven distribution of charge on a polar molecule. This makes one side more negative and one side more positive. Think about the tug of war. The atom with a stronger electronegativity will become slightly negative because the electrons are closer to the central atom. This is the opposite for the other atom.
Polarity: When a compound contains two opposite poles in which attract or propel one another. It is when a compound contains a negative and positive charge on either side of the molecule.
Polar: When there is an unevenly distributed charge creating an inbalance
Non-Polar: When there is an evenly distributed charge creating balance.
Finding Polarity: Within your compound, take the bigger EN and subtract the smaller EN to find the Polarity/electronegativity difference.
Electronegativity differences: Slightly polar = 0.0-0.5 Polar = 0.5-1.7 non-polar = 0.0-0.0 Ionic Bonding = 1.8-4.0
Where to find electronegativity (EN): Om the periodic table, the EN is found with every element.
Intra/Intermolecular Bonds/Forces(L6):
Intermolecular Bonds/Forces: Bonds between molecules caused by a force of attraction. These can only happen in molecular/covalent compounds. Generally, these bonds can be broken easily and are much weaker than intermolecular bonds.
London Dispersion Force (LDF): Weakest Intermolecular force. Created by the positive proton in nucleus and electrons from neighboring molecules. The only Intermolecular force which acts on non-polar molecules. Occurs In ever intermolecular bond.
Dipole Dipole Force (D-D): The second strongest intermolecular force. Occurs between polar molecules. These bonds occur because of attractions between a positive pole on one molecule and the negative pole on a neighboring molecule.
Hydrogen Bonds (H-Bonds): Occurs between molecules that contain hydrogen atoms directly attached to highly electronegative elements. This includes Nitrogen, Oxygen, and Florine. Because these elements have high electronegativity, they nearly strip the electrons away from the hydrogen making it become very positive thus very polar. Remember that hydrogen bonds only apply to molecules containing nitrogen, oxygen, or Florine with hydrogen bonded to it.
Intramolecular Bonds/Forces: Are bonds between atoms within a molecule. Includes covalent, ionic, metallic bonds. These are the strongest bonds.
Strongest to weakest bonds: Covalent, Ionic, Metallic, Hydrogen Bonds, D-D, LDF.
Poly atomic Ions and Metallic Bonding (L2):
Metallic Bonds: Occur when two or more nuclei and simultaneously attracted to the same electron. The valance electrons are shared between all atoms and are de localized meaning they move between atoms. Electrons are basically constantly moving and replacing each other.
Network Solids: Are created when millions of atoms of the same element or two different elements covalently bond to form patterned structures.
VSEPR and Stereochemistry (L3):
AXE: Where A represents the central atom, X represents bonded atom, and E represents LP. m is the number of bondeed atoms and n is the number of LP.
VSPER: The prediction of the 3-D shapes of a molecule.
Trigonal Planar: When 3 atoms bond to a central atom. 120 degrees. 0 LP
Linear: When bonded atoms create a straight line. 180 degrees. 0 LP
Tetrahedral: When 4 atoms bond to a central atom. 109.5 degrees. 0 LP
Trigonal Pyramidal: When 3 atoms bond to a central atom and but the central atom has an LP. 107.5 degrees. 1 LP
Bent V-shape: When two atoms bond to a central atom, but the central atom has two LP. 104.5 degrees. 2 LP
Lesson 5 Application of L 1-4/ Practice
Effects of intermolecular bonds on physical properties(L7):