2.2 Molecular Compounds pp. 61 – 69

First Some Useful Vocabulary  Diatomic molecules – consist of two atoms sharing a covalent bond  Polyatomic molecules – consist of more than two atoms sharing covalent bonds  Lone pair – A pair of valence electrons not involved in bonding  Bonding Capacity – the number of electrons lost, gained, or shared by an atom when it bonds (e.g. N has a bonding capacity of 3)

 Covalent Bond – the attractive force between two atoms that results when electrons are shared.  Octet Rule – when the shared electrons are included, each atom has 8 valence electrons (2 in the case of hydrogen).  Lewis Structure – a representation of the number, arrangement, and types of atoms in a molecule where dashes represent covalent bonds and all valence electrons are shown.  Structural Formula – similar to a Lewis structure but lone pairs of electrons are not shown.

Drawing Lewis Structures and Structural Formulae for Molecular Compounds 1. The element with the largest bonding capacity is placed in the middle 2. Place the other atoms around the central atom separating them as far as possible from eachother. 3. Add up the number of valence electrons available in each atom. Add or subtract electrons as necessary for a polyatomic ion. 4. Place a dashed line (representing two electrons) between the central atom and each of the surrounding atoms. 5. Place the remaining electrons as lone pairs around the peripheral atoms until the octet rule is satisfied.

5. If octets are not complete, move lone pairs into bonding position until all octets are complete. 6. If the peripheral atoms all have complete octets and there are pairs of electrons remaining, place them as lone pairs on the central atom. 7. To give the structural formula, remove the dots representing the lone pairs. 8. Place brackets around polyatomic ions and write the charge outside the brackets.

Lewis Structures & Structural Formulae  Draw the Lewis structure and structural formulae of the following: 1. Oxygen gas, O 2 2. Water, H 2 O 3. Carbonate ion, CO Sulfur trioxide, SO 3

Lewis Structures for Covalent Bonds  Draw the Lewis structure for two Fluorine atoms and their reaction forming the diatomic Fluorine Molecule.  Draw the structural formula of this molecule.

Polyatomic Ions & Coordinate Covalent Bonds  Polyatomic ions are covalently bonded groups of atoms with an overall charge  They sometimes act as the non-metal in forming an ionic bond with a metal (e.g. Sodium carbonate Na 2 CO 3 )  They sometimes act as the metal in forming an ionic bond with a non-metal (e.g. Ammonium chloride NH 4 Cl)  Coordinate Covalent Bonds are bonds in which both of the shared electrons come from the same atom. This can often form Polyatomic Ions  e.g. NH 4 formed from NH 3 + H +

Lewis Structure for Polyatomic Ion  Draw the Lewis structure for the formation of NH 4 +

The Strength of Covalent Bonds Covalent molecules that contain only a few atoms are called simple covalent structures. weak bonds between molecules strong bonds within molecules Most substances that contain simple covalent molecules have low melting and boiling points and are therefore liquids or gases at room temperature, e.g. water, oxygen, carbon dioxide, chlorine and hydrogen. Why? The covalent bonds within (INTRAMOLECULAR) these molecules are strong but the bonds between (INTERMOLECULAR) molecules are weak and easy to break.

What is the structure of a molecular compound/solid? A few substances that contain simple covalent molecules are solid at room temperature. These are molecular solids. The solid is formed because millions of iodine molecules are held together by weak forces of attraction to create a 3D molecular lattice. Two iodine atoms form a single covalent bond to become an iodine molecule. weak forces of attraction What properties would you expect molecular solids to have with this type of structure? Iodine is a molecular solid at room temperature.

What are the properties of molecular compounds? low melting and boiling points; usually soft and brittle – they shatter when hit. The properties of a molecular solid, such as iodine, are: Why do molecular solids have these properties? The weak forces of attraction between the molecules can be broken by a small amount of energy. This means that the molecular solids are soft and brittle and melt and boil at low temperatures. Molecular solids are also unable to conduct electricity because there are no free electrons or ions to carry a charge. cannot conduct electricity.

Homework  Read pp. 61 – 69  Answer the following questions:  p. 69 # 1 – 5