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Bonding - SCH3U

Molecular Bonding Unit

NameRelated Information
Drawing Lewis Structures 1)Write element symbol 2) Drow one dot for each valence electron 3) Dots are spread over the four sides
Determining the Number of Valence Electrons The number of valence electrons is equal to the group number with the exception of helium
Representation of Ionic compounds
Covalent bonding H-Cl (Cl would have 3 lone pairs)
Structural Formula H-Cl (No lone pairs are shown)
Double and Triple bonds Atoms can share 4 electrons for a double bond, or 6 electrons for a triple bond (ig. O=O , however each O would have 3 lone pairs)
Bonding Capacity The number of shared electron pairs, with covalent bonds, that an atom can form
Drawing Lewis Structures (First 3 steps) 1) Arrange the element symbols so the the element with the highest bonding capacity is the central atom. 2) Add up the number of valence electrons (ass 1 for each negative charge, and subtract on for each positive charge) 3) Draw a skeletal structure
Drawing Lewis Structures (Last 3 steps) 4) Complete the octets and peripheral atoms (Hydrogen won't have lone pairs) 5) Place extra electrons on the central atom 6) If the central atom doesn't have an octet, try forming multiple bonds by moving lone pairs
Co-ordinate Bovalent Bonds Both electrons are donated by the same atom
VSEPR Accronym Valence Shell Electron Pair Repulsion Thory
Covalent Bond Characteristics Highly directions, a molecule of a substance, and has a definite shape
Why can we Predict Covalent Bond Shapes We assume that the lone electron pairs in the valence shell of atoms stay as far apart as possible
VSEPR Formula: AX Name of Shape: Linear
VSEPR Formula: AX2 Name of Shape: Linear
VSEPR Formula: AX3 Name of Shape: Trigonal Planer
VSEPR Formula: AX4 Name of Shape: Tetrahedral
VSEPR Formula: AX3E Name of Shape: Trigonal Pyramidal
VSEPR Formula: AX2E2 Name of Shape: Bent
Intramolecular Force Attractive forces between atoms and ions within a molecule
Intermolecular Forces Attractive forces between molecules (ig. Vander Walls forces) Weaker that intramolecular forces
Proof of Intermolecular Forces If covalent bonds were the only forces at work, most molecular compounds would be gasses as the attraction would not be strong enough to group them as liquids or solid
Dipole - Dipole attraction Force of attraction between oppositely charges ends of polar molecules. Quantitative measure of the charge separation in a bond.
Dipole - Dipole arrows Electronegativities are used to determine the direction of the dipole moment, the arrow will point towards the more electronegative element
London Dispersion Attractive forces between all molecules, including non polar ones. Result of temporary displacement of electron cloud around atoms, temporary short lived dipoles, weaker than dipole - dipole
Hydrogen Bonding Strong dipole-dipole force between the positive hydrogen atom of one molecule and the highly electronegative atom of another molecule. (H w/ F,O or N created hydrogen bonding due to strong electronegativity)
Non-Polar Covalent Bonds 2 identical atoms, or atoms with similar electronegativities. Electronegativity difference must be between 0.4
Polar Covalent Bonds If electronegativities are unequal. the electrons spend more time with the more electronegative bond
Polar molecules Not all molecules with polar bonds are polar molecules A molecule can have polar bonds but be non polar if the molecular symmetry causes the dipoles to cancel eachother Totally symmetrical atoms are non polar
Heteronuclear diatomic molecules These molecules have two different atoms and is polar
Molecular Properties of Polar Compounds Significant intermolecular attraction (Hydrogen bonding and dipole-dipole) Higher melting and boiling points Dissolve in polar compounds but not non polar
Molecular Properties of Non-Polar Compounds Small intermolecular attraction (London forces) lower melting and boiling points Dissolve in non polar compounds but not polar
Electronegativity Differeces Non polar covalent (0-0.4) Polar covalent (0.41-1.66) ­Ionic (1.67 or greater)
Created by: JayMay3
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