1. Forming Molecular Bonds

2.  What is a covalent bond?  The chemical bond that results from the sharing of electrons  Non-metals combine to acquire a full valence shell of 8 valence electrons  Example: 7 valence e - 8 valence e -

3.  Degree of sharing Non-polar Covalent Polar Covalent Ionic e - are shared equally e - shared unequally one atom takes e - from another Na Cl e-e-

4.  Types of covalent bonds  Single Covalent bonds- (also called sigma bonds) When a single pair of electrons is shared  Ex: H H  Group 7A: will form single covalent bonds  Group 6A: will form two single covalent bonds Ex: H 2 O  Group 5A: will form three single covalent bonds Ex: NH 3 - Ammonia  Group 4A: will form four single covalent bonds Ex: CH 4 - methane

5.  Types of covalent bonds continued…  Multiple covalent bonds: Double or Triple bonds  Double covalent bond- when two pairs of electrons are shared.  Ex: O 2 (draw lewis structure)  Triple covalent bond- formed when three pairs of electrons are shared between two atoms. Ex: N 2 shares three pairs of electrons.  Pi bonds π - multiple bond consists of one sigma and one pi bond.  triple bond- one sigma and two pi bonds.  The shorter the bond the stronger the bond. Triple bonds are shorter.

6.  Rules for drawing dot structures: 1.Calculate the number of valence e - each atom contributes. Divide this number in half to get the number of pairs. : : : : : : ex: CBr 4 C = 4 Br = 7 x 4 = 28 32 e - 16 prs

7.  3. Use pairs of e - (as single covalent bonds) to attach all the other atoms to the center atom. ex: CBr 4 C Br 4. Put lone pairs of e - on the outside atoms until each atom has 8 electrons (4 pairs) or 1 pair on hydrogen. 5. Put any leftover pairs on the center atom so that it also has 4 prs around it. : : : : : : : : : : : :

8.  Lewis dot structures continued…  2. Decide which element will be the center at usually the one that has fewer atoms or the lower electronegativity CH 4 SO 2 PCl 5

9.  Lewis dot structures continued…  CH 4  What is the central atom? C How many valence electrons does it have? 4 How many hydrogens are there? 4 How many valence electrons do each have? 1 H H H H Now, join the electrons with a bond How many sigma bonds are there? 4

10.  diatomic elements H2H2 O2O2 Br 2 F2F2 I2I2 N2N2 Cl 2 atoms share electrons in order to have 8 valence e - (2 for hydrogen)

11.  Naming Covalent Compounds Covalent compounds are named by adding prefixes to the element names. ‘Covalent’ means both elements are nonmetals. A prefix is added to the name of the first element in the formula ONLY if more than one atom of it is present. A prefix is ALWAYS added to the name of the second element in the formula The second element will use the form of its name ending in ‘ide’.

12.  Naming Covalent Compounds Prefixes SubscriptPrefix 1mono- 2di- 3tri- 4tetra- 5penta- SubscriptPrefix 6hexa- 7hepta- 8octa- 9nona- 10deca- Note: When a prefix ending in ‘o’ or ‘a’ is added to ‘oxide’, the final vowel in the prefix is dropped.

13.  Naming Binary Covalent Compounds: Examples N 2 S 4 dinitrogen tetrasulfide NI 3 nitrogen triiodide XeF 6 xenon hexafluoride CCl 4 carbon tetrachloride P 2 O 5 diphosphorus pentoxide SO 3 sulfur trioxide 1mono 2di 3tri 4tetra 5penta 6hexa 7hepta 8octa 9nona 10deca * Second element ends ‘ ide ’ * Drop –a & -o before ‘oxide’

14.  Writing Chemical Formulas: A Review I. Ionic Compounds II. Covalent Compounds

15.  Writing Formulas for Covalent Compounds The names of covalent compounds contain prefixes that indicate the number of atoms of each element present. Remember:  Binary compounds contain only two elements, both of which are nonmetals  When in covalent compounds  atoms DO NOT have charges  Subscripts are determined directly from the prefixes in the name. First element : if there is only one atom of that element in the formula (its subscript will be 1) and there should NOT be a prefix Second element: will ALWAYS have a prefix will ALWAYS end in -ide

16.  Writing Formulas for Binary Covalent Compounds: Examples nitrogen dioxide NO 2 diphosphorus pentoxide P2O5P2O5 xenon tetrafluoride XeF 4 sulfur hexafluoride SF 6 1mono 2di 3tri 4tetra 5penta 6hexa 7hepta 8octa 9nona 10deca * Second element ends in ‘ide * Drop –a & -o before ‘oxide’

17.  Writing Formulas: Practice carbon tetrafluorideCF 4 Na 3 PO 4 sodium phosphate Cu 2 SO 4 copper (I) sulfate AnalysisIf “Yes” The compound is covalent: the prefixes give the subscripts. * Are there prefixes present The compound is ionic: subscripts must be determined by balancing charges prefixes  covalent  prefixes indicate subscripts metal  ionic  balance charges  3 Na 1+ needed for 1 PO 4 3- metal present  ionic  balance charges  2 Cu 1+ needed for 1 SO 4 2- Al 2 S 3 aluminum sulfide metal present  ionic  balance charges  2 Al 3+ needed for 3 S 2- N 2 O 5 dinitrogen pentoxide prefixes  covalent  prefixes indicate subscripts NH 4 NO 3 ammonium nitrate polyatomic ion present  ionic  balance charges  1 NH 4 1+ needed for 1 NO 3 1- PbO 2 lead (IV) oxide metal present  ionic  balance charges  1 Pb 4+ needed for 2 O 2- Fe 2 (CO 3 ) 3 iron (III) carbonate metal present  ionic  balance charges  2 Fe 3+ needed for 3 CO 3 2-