1. Molecular Structure and Shape
Covalent Compounds
Molecular Structure and Shape

2. Lewis Structures Covalent compounds share electrons.
Goal is to achieve “Noble Gas” configuration.
Hydrogen wants to be like Helium (2 electrons)
All other nonmetals want to have 8 electrons. This is known as the octet rule.
Atoms achieve an octet by sharing 1, 2, or 3 electrons with another atom.
Single bond – each atom contributes ONE electron for a pair
Double bond – each atom contributes TWO electrons for two pairs
Triple bond – each atom contributes THREE electrons for three pairs

3. Lewis Structures
Atoms can also share electrons with more than one atom.
Ex. An oxygen atom shares 2 electrons with 2 hydrogen atoms in the formation of water. The oxygen shares 1 electron with each hydrogen atom, forming 2 single bonds.
Lewis structures show the electrons that can be shared and therefore can be used to figure out how atoms combine in covalent compounds.

4. Lewis Structures
Remember, a Lewis Structure (or electron dot diagram) is drawn by writing the symbol for an element and then surrounding it with dots that represent the valence electrons.
Example:
O H

5. Lewis Structures
Electrons are then SHARED between atoms so they can achieve their Noble Gas configuration.
Example:
H O H

6. H O H H O H Structural Formulas
Lewis Structures are converted to structural formulas by drawing a line in place of a SHARED pair of electrons.
H O H
H O H

7. Draw the Lewis Structure for the following elements.
Hydrogen
Carbon
Oxygen
Nitrogen
Sulfur
Chlorine
Fluorine
Boron
Selenium

8. Identify the number of electrons needed to complete the octet
Identify the number of electrons needed to complete the octet. -This is how many electrons will be shared.
Hydrogen
Carbon
Oxygen
Nitrogen
Sulfur
Chlorine
Fluorine
Boron
Selenium

9. Using the element Lewis Structure, draw the Lewis Structure for the following compounds.
Hydrogen, H2
Chlorine, Cl2
Oxygen, O2
Nitrogen, N2
Methane, CH4
Ammonia, NH3
Hydrofluoric acid, HF
Carbon dioxide, CO2
Boron trichloride, BCl3
Ethane, C2H6

10. Molecular Shapes
Shapes are determined by the number of shared and unshared pairs of electrons that surround the central atom.
Molecules arrange themselves so that their electron pairs are as far apart as possible.
This is the VSEPR theory – Valence Shell Electron Pair Repulsion theory

11. Shapes you need to know Linear Tetrahedral, Trigonal Pyramid, Bent
Steric Number 2 or higher
Tetrahedral, Trigonal Pyramid, Bent
Steric Number of 4
Trigonal Planar
Steric Number of 3
Octahedral
Steric Number of 6

12. Sigma and Pi Bonds
Sigma Bonds are the first bond that forms between two atoms.
Sigma bonds are the only bonds that are used to determine the steric number and molecular shape.
Sigma bonds form in the space directly between two atomic nuclei.
Pi Bonds are the second or third bond that forms between the same two atoms.
Pi bonds are NOT used to determine molecular shape. However, they do shorten the length of the bond – pulling the nuclei closer together.
Pi bonds form around the sigma bond – either top and bottom or front and back.

13. Sigma and Pi Bonds

14. Sigma and Pi Bonds Single bond Double bond Triple bond 1 sigma bond
Ethane, C2H6
Double bond
1 pi bond
Ethene, C2H4
Triple bond
2 pi bonds
Ethyne, C2H2

15. Polarity Bonds and Molecules are either polar or nonpolar.
Polar Bonds form when there is unequal sharing of the electrons.
This is due to a difference in electronegativity.
The more electronegative nucleus gets “a bigger share”.
Polar Molecules form when the polar bonds and unshared pairs surrounding the central atom are asymmetrical in their 3D arrangement.
Certain shapes will usually be polar.
Trigonal pyramid, bent
All shapes can be polar.

16. Intermolecular Forces
Forces BETWEEN molecules
Hydrogen Bonds occur when a hydrogen is covalently bonded to a very electronegative atom (F, Cl, O)
Strongest of intermolecular forces
Dipole Interactions occur with polar molecules. The positive end of one molecule is attracted to the negative end of another molecule.
Dispersion forces occur as a result of the motion of electrons within atoms.
Weakest of intermolecular forces

17. Review of Compounds
Ionic, Covalent, Metallic

18. Ions and Ionic Compounds
Remember an ion is an atom that has lost or gained electrons
Cations – positive – lost electrons
Anions – negative – gained electrons
Ionic Compounds form when 2 or more atoms are joined by the loss and gain of electrons
ALWAYS cation + anion
Cation is always first and anion is always second.

19. Ionic Bonds Cation + Anion
Ions are joined by the transfer of electrons
Creates Electrostatic forces (attraction of opposite charges) that hold the ions together
Ionic Compounds are composed of a continuous arrangement of oppositely charged ions.
NOT a single separate unit

20. Ionic Bonds
Bonds between atoms will be ionic when there is a LARGE difference in electronegativity between the atoms.
> 1.7 ∆EN
Metal + Nonmetal
Opposite sides of the periodic table – large differences in electronegativity

21. Properties of Ionic Compounds
Metal + Nonmetal (usually)
Crystalline SOLIDS at room temp. (most)
Crystal – repeating geometric pattern
Brittle
HIGH melting points
HIGH boiling points
Some are so high it takes extreme conditions to get them to change to gas
Conduct electrical currents when melted or dissolved in water

22. Molecular Compounds Compounds that form when atoms SHARE electrons
Forms a covalent bond
NEVER contain ions
NEVER have charges
Contain only nonmetals

23. Properties of Covalent Compounds
Nonmetal + Nonmetal
Can be solids, liquids, or gases at room temperature
State is determined by the bond strength (compounds with stronger bonds tend to be solids)
LOW melting points
LOW boiling points
POOR conductors of electricity under any conditions
Can be Polar or Nonpolar
Depends on how the electrons are shared
Polar compounds are better conductors

24. Metallic Compounds Metal + Metal
Consists of positive metal ions with a “sea of electrons”
Electrons are free floating – are not attached to any one atom or ion
HIGH melting points
HIGH boiling points
HIGH conductivity
Malleable, Ductile, High luster (shine)
All properties are a result of the “free” electrons.