Covalent Bonds

1. Covalent Bonds The joy of sharing!

2. Covalent Bonds • Covalent bonds: occur between two or more nonmetals; electrons are shared not transferred (as in ionic bonds) • The result of sharing electrons is that atoms attain a more stable electron configuration.

3. Covalent Bonds • Most covalent bonds involve: • 2 electrons (single covalent bond), • 4 electrons (double covalent bond, or • 6 electrons (triple covalent bond).

4. Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons) • H2 HBr • CCl4 O2 • N2 CO

5. Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons) • H2 HBr • CCl4 O2 • N2 CO

6. Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons) • H2 HBr • CCl4 O2 • N2 CO

7. Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons) • H2 HBr • CCl4 O2 • N2 CO

8. Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons) • H2 HBr • CCl4 O2 • N2 CO

9. Lewis structures (electron dot structures) show the structure of molecules. (Bonds can be shown with dots for electrons, or with dashes: 1 dash = 2 electrons) • H2 HBr • CCl4 O2 • N2 CO

10. Octet Rule • Octet Rule: The representative elements achieve noble gas configurations (8 electrons) by sharing electrons.

11. Writing Lewis Structures • Select a skeleton for the molecule (the least electronegative element is usually the central element). • Calculate N (the # of valence e- need by all atoms in the molecule of polyatomic ion. • Calculate A (the # of electrons available). • Calculate S (the # of electrons shared in the molecule) S = N – A • Place the S electrons as shared pairs in the skeleton. • Place the additional electrons as unshared pairs to fill the octet of every representative elements (except hydrogen!).

12. CO2- OH- NO3- SO42- CBr3- N22- CO32- NH4+ Lewis Structure Examples:

13. CO2- OH- NO3- SO42- CBr3- N22- CO32- NH4+ Lewis Structure Examples:

14. CO2- OH- NO3- SO42- CBr3- N22- CO32- NH4+ Lewis Structure Examples:

15. CO2- OH- NO3- SO42- CBr3- N22- CO32- NH4+ Lewis Structure Examples:

16. CO2- OH- NO3- SO42- CBr3- N22- CO32- NH4+ Lewis Structure Examples:

17. CO2- OH- NO3- SO42- CBr3- N22- CO32- NH4+ Lewis Structure Examples:

18. CO2- OH- NO3- SO42- CBr3- N22- CO32- NH4+ Lewis Structure Examples:

19. CO2- OH- NO3- SO42- CBr3- N22- CO32- NH4+ Lewis Structure Examples:

20. Electronegativity • We’ve learned how valence electrons are shared to form covalent bonds between elements. So far, we have considered the electrons to be shared equally. However, in most cases, electrons are not shared equally because of a property called electronegativity.

21. Electronegativity • The ELECTRONEGATIVITY of an element is: the tendency for an atom to attract electrons to itself when it is chemically combined with another element. • The result: a “tug-of-war” between the nuclei of the atoms.

22. Electronegativity • Electronegativities are given numerical values (the most electronegative element has the highest value; the least electronegative element has the lowest value) • **See Figure 6-18 p. 169 (Honors) • Most electronegative element: Fluorine (3.98) • Least electronegative elements: Fr (0.70), Cs (0.79)

23. Electronegativity • Notice the periodic trend: • As we move from left to right across a row, electronegativity increases (metals have low values nonmetals have high values – excluding noble gases) • As we move down a column, electronegativity decreases. • The higher the electronegativity value, the greater the ability to attract electrons to itself.

24. Nonpolar Bonds • When the atoms in a molecule are the same, the bonding electrons are shared equally. • Result: a nonpolar covalent bond • Examples: O2, F2, H2, N2, Cl2

25. Polar Bonds • When 2 different atoms are joined by a covalent bond, and the bonding electrons are shared unequally, the bond is a polar covalent bond, or POLAR BOND. • The atom with the stronger electron attraction (the more electronegative element) acquires a slightly negative charge. • The less electronegative atom acquires a slightly positive charge.

26. Polar Bonds • Example: HCl • Electronegativities: • H = 2.20 • Cl = 3.16 - + H Cl

27. Polar Bonds • Example: H2O • Electronegativities: • H = 2.20 • O = 3.44

28. Polar Bonds • Example: H2O • Electronegativities: • H = 2.20 • O = 3.44

29. Polar Bonds • Example: H2O • Electronegativities: • H = 2.20 • O = 3.44

30. Predicting Bond Types • Electronegativities help us predict the type of bond: covalent (nonpolar) covalent (slightly polar) covalent (very polar) ionic

31. Polar Molecules • A polar bond in a molecule can make the entire molecule polar • A molecule that has 2 poles (charged regions), like H-Cl, is called a dipolar molecule, or dipole.

32. Polar Molecules • The effect of polar bonds on the polarity of a molecule depends on the shape of the molecule. • Example: CO2 O = C = O shape: linear *The bond polarities cancel because they are in opposite directions; CO2 is a nonpolar molecule.

33. Polar Molecules • The effect of polar bonds on the polarity of a molecule depends on the shape of the molecule. • Water, H2O, also has 2 polar bonds: • But, the molecule is bent, so the bonds do not cancel. • H2O is a polar molecule.