LEWIS STRUCTURES AND VSEPR THEORYmrjvacosta.weebly.com/uploads/1/4/2/6/14262365/...Octet rule – 2....
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LEWIS STRUCTURES AND VSEPR THEORY
Molecular Compounds
Molecule - a neutral group of atoms joined together by covalent bonds
Diatomic molecule – molecule consisting of two atoms
Molecular Compounds
Molecular Formula – Shows how many atoms of each element a molecule contains Chemical formula of a covalent compound
Doesn’t tell you anything about the molecule’s structure,
arrangement or shape
Molecular Structural Ball-and-Stick model Space-filling model Formula formula
Types of Bonds
Single bond – two atoms share 1 pair of electrons
Double Bond – two atoms share 2 pairs of electrons
Triple Bond – two atoms share 3 pairs of electrons
Single Covalent Bonds
Two atoms held together by sharing a pair of electrons are joined by a single covalent bond.
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Single Covalent Bonds
The halogens form single covalent bonds in their diatomic molecules. Fluorine is one example.
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Single Covalent Bonds
The hydrogen and oxygen atoms attain noble-gas configurations by sharing electrons.
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Single Covalent Bonds
The ammonia molecule has one unshared pair of electrons.
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Double and Triple Covalent Bonds
Each nitrogen atom has one unshared pair of electrons.
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Double and Triple Covalent Bonds
Carbon dioxide is an example of a triatomic molecule.
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Coordinate Covalent Bonds
A covalent bond in which one atom contributes both bonding electrons
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Lewis Structures
Representation of a molecule that depicts how the valence electrons are arranged among the atoms in the molecule Bonding pairs – pairs of electrons shared between two atoms Lone pairs- pairs of electrons that are not involved in bonding
Octet rule – 2nd row elements typically obey the octet rule when forming covalent compounds
Duet rule – hydrogen
Lewis Structures
Determine the total # of valence electrons in atoms to be combined
Draw the skeleton of the structure and connect atoms with single bonds Carbon is usually the
central atom Otherwise the least
electronegative atom Hydrogen is NEVER the
central atom
Step 1 Step 2
Lewis Structures
Add electrons until each atom has a complete octet of electrons.
Count to see if the total # of electrons equals the amount calculated in step one. If yes, structure is correct If no, move lone pairs in
order to make double or triple bonds
Step 3 Step 4
Exceptions to the Octet Rule
Expanded Octets
Incomplete Octets
Bond Polarity
Polar Covalent Bond – covalent bond in which one electrons are shared unequally More electronegative atom attracts electrons more
strongly and gains a partial negative charge
Less electronegative atom has a slightly positive charge
Bond Polarity
Electronegativity Difference
Bond Type Element Type Example
0.0 – 0.4 Non-polar covalent
Non-metal + non-metal
0.4 – 2.0
Polar covalent Non-metal + non-metal
> 2.0 Ionic (polar) Metal + non-metal
What is the name of 007's Eskimo cousin?
Polar Bond
Valence Shell Electron Pair Repulsion
Structure or shape of molecule is determined by minimizing the repulsions between electron pairs Bonding and lone pairs are positioned as far apart as
geometrically possible
Polarity
Polar Molecule – a molecule with a distinct positive and distinct negative end Must have polar bonds within the molecule Must have correct geometry
Intermolecular Forces
Ionic Bonding – results from the electrostatic attraction between positive and negative ions Very strong force that results in room temperature solids
with high melting and boiling points
Intermolecular Forces
Dipole-Dipole - Electrostatic attraction between a partially negative end of one molecule and a partially positive end of another molecule Occurs between polar molecules
Intermolecular Forces
Hydrogen Bonding - An especially strong dipole-dipole that involves molecules with highly electronegative atoms bonded to hydrogen Usually N, O, or F
Intermolecular Forces
London Dispersion Forces Result from instantaneous, non-permanent dipoles
created by random electron motion Temporary dipoles cause weak and temporary
electrostatic attraction between molecules Interactions between non-polar molecules Molecules with more mass and thus electrons have stronger
LDF
Intermolecular Forces
Ion-dipole – electrostatic attraction between an ion and a polar molecule
Ion-induced dipole – electrostatic attraction between an ion and a non-polar molecule
Dipole-induced dipole – electrostatic attraction between a polar and non-polar molecule
Affect of IMF on Chemical Properties
Melting Point and Boiling Point – the stronger the intermolecular force the higher the melting or boiling point
Solubility – like dissolves like
Surface Tension – adhesion of molecules at the surface of a liquid
Why did the white bear dissolve in water?
Because it was polar