Chapter 9: Covalent Bonding

1. Chapter 9: Covalent Bonding

2. Review • Noble gases are the most stable • Have full outer energy level • Do not react with other elements to form bond • Metals and Nonmetals react • form binary ionic compounds • electrons are transferred • attain a noble gas configuration

3. Sometimes two atoms that BOTH need to gain electrons to become stable will have a similar attraction for electrons • These atoms SHARE electrons • RECALL: Octet rule states that atoms will gain, lose, or share electrons in order to become stable

4. What is a covalent bond? • Covalent bond- results from the sharing of valence electrons • The shared electrons are considered to be part of the complete outer energy level • Forms between two non-metals

5. Molecules • Forms when two or more atoms bond covalently Ex. Carbohydrates, proteins, DNA, and fats

6. Formation of a Covalent Bond • Hydrogen (H2), Nitrogen (N2), Oxygen (O2), Fluorine (F2), Chlorine (Cl2), Bromine (Br2), and Iodine (I2) are all diatomic molecules • Diatomic molecules form because they are more stable than the individual atoms • Consider two individual atoms of Fluorine...

7. Fluorine’s electron configuration is as follows: 1s22s22p5 • Each fluorine atom has 7 valence electrons • As two atoms of Fluorine approach each other, a repulsive force occurs • The electrons repulse each other but there is also an attraction between the nuclei of each atom and the electrons • A maximum point of attraction occurs and the force of attraction balance out the forces of repulsion (most stable) • At this point, the covalent bond forms

8. Covalent Vocabulary • Lone pairs-the unshared pair of electrons • Bonding pair- the shared pair of electrons Ex. Look at the dot structure of F2

9. Single Covalent Bonds • Formed by the sharing of one pair of electrons • Hydrogen and the halogen group elements will form single covalent bonds with an identical atom • Group 6A elements will form two single covalent bonds (ex. H2O, Water) • Group 5A will form three single covalent bonds (ex. NH3, Ammonia)

10. Page 244, Practice Problems Draw the Dot Structures for these molecules: • PH3 • H2S • HCl • CCl4 • SiH4

11. Sigma Bonds • Sigma bonds are single covalent bonds • Represented by the Greek letter sigma (σ) • Occurs when electron pair is shared in an area centered between two atoms and there is an overlap in the valence orbital

12. Multiple Covalent Bonds • Double bonds- two pairs of electrons are shared (ex. O2) • Triple bonds – three pairs of electrons are shared (ex. N2) • Multiple bonds consist of a sigma bond and at least one pi bond (π) • Double bond has one sigma and one pi bond • Triple bond has one sigma and two pi bonds

13. Strength of Covalent Bonds • Strength depends on how much distance separates the nuclei • Bond Dissociation energy • Extra Credit Assignment – Due Dec. 1 • READ PAGES 246-247 • DO page 247 6-11 (6) • READ PAGES 252-258 • DO page 258 42-47(6) • READ PAGES 259 – 262 • DO page 259 54-58 (5)

14. 9.2 Naming Binary Molecules • First element is always named first, using entire elements name • The second element is named with the ending –ide • Prefixes are added to each name to indicate the number of atoms present • Exception: if there is only one of the first element omit prefix

16. Page 249, 13-17 13. CCl4 14. As2O3 15.CO 16.SO2 17.NF3

17. Page 249, Table 9-2

18. Naming Acids • Water solutions of some molecules are acidic (ACIDS) • If in solution, the compound produces H+ ions, then it is acidic • Example: HCl hydrochloric acid • Two Types of Acids: • 1. Binary Acids • 2. Oxyacids

19. Binary Acids • Contains hydrogen and one other element • When naming, use prefix hydro- to indicate the presence of hydrogen • Add the ending –ic, and then write the word acid at the end of the name. • Ex. HBr hydrobromicacid

20. Oxyacids • Acids containing an oxyanion such as SO4 • In naming, do not use the hydro prefix, you will need to change the suffix on the anion • If the ending is –ate change it to –ic • If the ending is –ite change it to –ous HNO3 nitric acid HNO2 nitrous acid

21. Page 250, 18-22 18. HI 19. HClO3 20. HClO2 21. H2SO3 22. H2S

22. 9.3 Molecular Structures • Structural formulas use the letter symbols and the dots to illustrate relative positions of atoms • The central atom is the one with the least attractive force • Hydrogen is always terminal atom due to only sharing one possible electron Ex. NH3

23. So, How do polyatomic ions form anyway? • Example: ClO4-1 PO4-3 NH4+1

24. Resonance Structures • It is possible to have more than one correct Lewis Structure • These are called resonance structures • Ex. Nitrite ion NO2-1 • These are common molecules with more than one structure: ozone (O3), nitrate, nitrite, sulfite, and carbonate

25. 9.4 Molecular Shape • Molecular geometry is important ot understanidng the stregth of bonds • VSEPR Model determines shape • Valence Shell Electron Pair Repulsion • Based on minimizing repulsion of shared and unshared pairs of electrons • Memorize shapes on page 260

26. 9.5 Electronegativity and Polarity • Polar bonds occur when electrons are not shared equally, resulting in unequal distribution of charge and the formation of a dipole • spatial arrangement of polar bonds in a molecule determines the overall polarity of a molecule