Read Sections 4.6, 4.7, 4.8, 4.9 and 4.10 before viewing the slide show.

Unit 15 Molecular Compounds Lewis Structures for Molecular Compounds and Ions Polyatomic Ions Naming Molecular Compounds

Lewis Structures for Molecular Compounds and Ions (4.7, 4.10) Nonmetals – those to the right of the stair-step line can react with each other by sharing electrons The driving force is similar to that for ionic compounds – the attainment of an octet of electrons as summarized in the Octet Rule In molecular compounds, though, the atoms do not form individual ions, they share electrons throughout their structure In this unit, we will look at drawing Lewis structures representing these molecular compounds as well as other aspects of these shared electron systems

Key Aspects of Lewis Structures (4.10) Key elements of a Lewis structure for a molecular compound include: The number of valence electrons drawn equals the total number of valence electrons in the compound or ion. Electrons are shared in such a way that each atom is effectively surrounded by eight electrons (two for hydrogen, other exceptions exist). All or some of the electrons around an atom may be shared with other atoms in the structure

An Approach to Drawing Lewis Structures (4.10) Several approaches exist for drawing Lewis structures – this is just one you might consider: 1.Draw skeletal structure for the atoms. Things that might be helpful are that hydrogen can only take one bond so it has to be on the outside; carbon typically has four bonds, nitrogen has three bonds, and oxygen has two bonds; polyatomic molecules and ions typically have a central atom surrounded by other atoms. 2.Connect each atom to its neighbor(s) with a single bond which consists of two electrons. The bond may be characterized by two electrons (:) or a single line (—) attaching the two atoms. 3.Place electrons around the outside of the atoms so each has eight total electrons including all of those bonded to it as well as the nonbonded (or lone pair) electrons. (Continued on Next Slide)

An Approach to Drawing Lewis Structures Continued (4.10) 4.Count the number of valence electrons the structure should have. This is the sum of the valence electrons on all atoms. For ions, the number of valence electrons is increased by the charge and for negative ions and reduced by the charge for positive ions. 5.Count the number of valence electrons in your structure. 6.If the valence electron counts match from steps 4 and 5 you are done. If there are too many valence electrons in your structure, take away a nonbonded pair and slide a nonbonded pair from an adjacent atom to make a double chemical bond. Continue this process until the number of valence electrons in the structure matches the number that should be present.

Examples of Drawing Lewis Structures (4.10)

Polyatomic Ions (4.9) An ion formed from two or more bonded atoms is called a polyatomic ion. These ions will often react as units and typically do not change during the course of a chemical reaction. There is not a wonderfully logical way of determining the names of polyatomic ions, though there are some patterns. Table 4.4 (on page 106 or the next slide) gives a short list of common polyatomic ions. Don’t panic – you don’t have to know all of these. For our purposes, the focus will be on five polyatomic ions: ammonium NH 4 + hydroxide OH - nitrate NO 3 - sulfate SO 4 2- phosphate PO 4 3- On the proctored midterm, I will not give you the information above – you will be expected to know these five. If you need others, I will give you this information for them.

Writing Formulas for and Naming Compounds with Polyatomic Ions (4.9) Writing formulas using polyatomic ions works just like it did with binary ionic compounds – think of the polyatomic ion as one big ion. In naming the compounds, the polyatomic ion retains its name – no changing of the ending or anything like that. Examples: Na + and SO 4 2- make Na 2 SO 4 called sodium sulfate NH 4 + and S 2- make (NH 4 ) 2 S called ammonium sulfide Ba 2+ and PO 4 3- make Ba 3 (PO 4 ) 2 called barium phosphate

Naming Binary Molecular Compounds ( ) Writing formulas for binary molecular compounds is a little trickier than for ionic compounds since the atoms don’t really have a charge (they are sharing electrons) and many different compounds can be formed from the same two elements. The naming is fairly direct. Prefixes are used (see the table to the right) to indicate the number of atoms of each type in the molecule. The mono- prefix is only used in reference to the second atom in the compound. Examples: CO carbon monoxide CO 2 carbon dioxide N 2 O dinitrogen monoxide (also known as nitrous oxide or laughing gas) NO 2 nitrogen dioxide N 2 O 4 nitrogen tetroxide N 2 O 5 dinitrogen pentoxide # of atomsPrefix 1mono- 2di- 3tri- 4tetra- 5penta- 6hexa- 7hepta- 8octa- 9nona- 10deca-

How many different elements (not atoms – elements) are in the compound? What types of elements are they? Metal keeps its name. For Groups 1 and 2, Al, Zn, and Ag no modification is necessary. For other metals, use a Roman numeral in parentheses to indicate the charge on each ion of the metal. Polyatomic ions retain their names. (Ionic compound) Metal keeps its name. For Groups 1 and 2, Al, Zn, and Ag no modification is necessary. For other metals, use a Roman numeral in parentheses to indicate the charge on each ion of the metal. Nonmetal changes its ending to –ide. (Molecular or covalent compound) First element retains its name. Second element switches to –ide ending. Each name is preceded by a prefix indicating the number of its atoms (di-, tri-, tetra-, etc.). Mono- is only used for the second element. Exactly twoMore than two Metal-nonmetal Nonmetal-nonmetal GSB Modified 2/1/2010 Summary of Inorganic Nomenclature