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