Chapter 5

1. Chapter 5 Molecular Compounds

2. Covalent Bonds • Covalent Molecule – A substance comprised of only nonmetals and metalloids • Covalent Bond • Formed by the sharing of electrons • each atom within the bond achieves an octet (8) except hydrogen (2) and boron (6) • Minimize the repulsions and maximize the attractions

3. Covalent Bonds • Diatomic molecules • Br I N Cl H O F

4. Naming Binary Molecular Compounds • The atoms of a molecular formula are written in the order of the elements electronegativity • Least electronegative atom will always be listed first. • Electronegativity – indicates how well an elements nuclei attract the electrons in a covalent bond

5. The Periodic Table and Electronegativity

6. Naming Covalent molecules • Translate the elemental symbol into the elements name • 1st element – give its name • 2nd element – gets ending –ide same as found in ionic compounds • Because electrons are shared • Indicate how many of each atom are present using the prefix system • All elements will receive a prefix except if there is only 1 of the first element

7. Naming Binary Molecular Componds The prefix mono- is omitted for the first element.

8. Problem • What is the name of PCl3?  • a.  phosphorus chloride • b.  phosphoric chloride • c.  phosphorus trichlorate • d.  trichlorophosphide • e.  phosphorus trichloride

9. Problem • The compound, P4S10, is used in the manufacture of safety matches. What is its name?  • a.  phosphorus sulfide • b.  phosphoric sulfide • c.  phosphorus decasulfide • d.  tetraphosphorus decasulfide • e.  phosphorus sulfide

10. Molecular Formulas • Translate the words

11. Covalent Bonds and the Periodic Table • Covalent bonds form between nonmetals • Achieve an octet (noble gas configuration) • Using electron dot symbols build • H2, H2O, CH4, O2, N2, HCN, CO2 • Least electronegative atom is central (except H)

12. Covalent Bonds and the Periodic Table • Elements follow octet rule • H gets 2 electrons – 1 bond • B has 3 valence electrons – can only form 3 bonds • Period 3 elements and higher can form more than an octet because of the empty d-orbitals

13. Covalent Bonds and the Periodic Table Number of bonds formed to achieve octet. Numbers in parentheses indicate possible numbers of bonds that result in exceptions to the octet rule.

14. Multiple Covalent Bonds • Single bond: A covalent bond formed by sharing one electron pair. • Double bond: A covalent bond formed by sharing two electron pairs. • Triple bond: A covalent bond formed by sharing three electron pairs.

15. Molecular Formulas and Lewis Structures • Molecular formula: A formula that shows the numbers and kinds of atoms in one molecule of a compound. • Structural formula: A molecular representation that shows the connections among atoms by using lines to represent covalent bonds. • Lewis structure: A molecular representation that shows both the connections among atoms and the locations of lone-pair valence electrons.

16. Molecular Formulas and Lewis Structures • The oxygen atom in H2O • shares 2 bonding pairs of electrons with two hydrogen atoms • Has 2 other pairs of valence electrons that are not shared in bonds – lone pairs

17. Drawing Lewis Structures • Sum up the valence electrons for all atoms in the molecule • Identify the atom in the lowest group # (except H) – Least electronegative atom • If both atoms in same group # identify the one in the highest numbered period. • Write this symbol as the central atom surrounded by all other atoms

18. Drawing Lewis Structures • Place a single bond (-, representing two electrons) between all external atoms and the central atom • Add up the electrons in the bonds and subtract this from the total valence electrons

19. Drawing Lewis Structures • Place all other electrons around the outer atoms (in pairs) first giving them an octet (except H) • Most electronegative atom first • Place all remaining electrons around the central atom as lone pairs • Check to see if all atoms have an octet • Yes – you are done • No – make multiple bonds

20. Drawing Lewis Structures • H, C, N, O, and halogen atoms usually maintain consistent bonding patterns: • H forms one covalent bond. • C forms four covalent bonds. • N forms three covalent bonds and has one lone pair of electrons. • O forms two covalent bonds and has two lone pairs of electrons. • Halogens form one covalent bond and have three lone pairs of electrons.

21. The Shapes of the Molecules • Molecular shapes can be predicted by noting how many bonds and electron pairs surround individual atoms and applying what is called the valence-shell electron-pair repulsion (VSEPR)model. • VSEPR model - the negatively charged clouds of electrons in bonds and lone pairs repel each other, and keep as far apart as possible.

22. The Shapes of the Molecules There are three step to applying the VSEPR model: • Step 1: Draw a Lewis structure of the molecule, and identify the atom whose geometry is of interest. • Step 2: Count the number of electron charge clouds surrounding the atom of interest. • Step 3: Predict molecular shape by assuming that the charge clouds orient in space so that they are as far away from one another as possible.

24. The Shapes of the Molecules • Linear molecules have bond anglesof 180°. • Planar triangular molecules have bond angles of 120°. • Tetrahedral molecules have bond anglesof 109.5°.

25. Polar Covalent Bonds and Electronegativity • As a rule of thumb • electronegativity differences of less than 0.5 - nonpolar covalent bonds • differences up to 1.9 indicate polar covalent bonds • differences of 2 or more indicate ionic bonds.

26. Problem • Select the most polar bond amongst the following.  • a.  C-O • b.  Si-F • c.  Cl-F • d.  C-F • e.  C-I

27. Polar Molecules • If the bonds in a molecule have a polarity then depending on the shape of the molecule, the molecule itself can be polar • Guaranteed nonpolar molecules • Linear, trigonal planar and tetrahedral if all peripheral atoms are identical

28. Characteristics of Molecular Compounds