1. Unit II: Bonding

2. Electron Pairs Recall that in the Lewis/electron dot diagram, only valence electrons are shown. We consider the shell that they are in to be the valence shell. These valence electrons give the element its chemical properties, and therefore determine its ability to form chemical bonds.

3. Electron Pairs This structure also tells us which electrons are paired and unpaired. Paired valence electrons are more stable & less likely to bond, so they’re called nonbonding pairs.

4. Electron Pairs Unpaired valence electrons strongly tend to bond by pairing up with the electron of another atom.

5. Electron Pairs These chemical bonds will happen in one of the following ways: 1. One atom gives an unpaired electron to another atom that also has an unpaired valence electron. 2. Two atoms each have an unpaired electron, so they share their valence electrons to form a single shared valence pair.

6. Symbols for Ions A superscript is used to show the magnitude and charge of an ion. – If a positive ion is formed from a sodium atom it is written as Na¹ +, the 1 can be omitted, so it means the same as Na + – If a negative ion is formed from the fluorine atom it is written as F -

7. Symbols for Ions – If an atom loses or gains more than one electron the charge number is given. A calcium atom loses two electrons - Ca 2+ An oxygen atom gains two electrons - O 2-

8. Symbols for Ions Rule: Atoms tend to lose or gain electrons so that they end up with an outermost occupied shell that is filled to capacity. See Figures on page 391.

9. Symbols for Ions If an atom has a small amount of valence electrons, it is more likely to lose those electrons. Therefore it forms a positive ion. If an atom has a valence shell that is almost filled, it is more likely to gain electrons to fill its valence shell. Therefore it forms a negative ion.

10. Symbols for Ions Write the symbol for the types of ions atoms tend to form above each group

11. Ionics Bonds An ionic bond results from a transfer of valence electrons. An electron transfer occurs when an atom that tends to lose an electron comes in contact with an atom that tends to gain an electron.

12. Ionic Bonds The result is that one atom ends up with a + charge, and the other atom ends up with a – charge. The combination of sodium and chlorine is the classic example.

13. Ionic Bonds

14. The two oppositely charged ions are now attracted to each other by electric force. This electric force of attraction is called an ionic bond.

15. Ionic Bonds All chemical compounds containing ions are referred to as ionic compounds. The compounds are totally different from the elements that make them. The charges in these compounds must always balance. See figure 23.9 & 23.10 on page 395.

16. Ionic Bonds Ionic compounds tend to be grouped together in a highly ordered three-dimensional array, often forming a crystalline structure. Ionic compounds tend to form from elements on opposite sides of the periodic table.

17. Covalent Bonds Covalent bonds result from a sharing of valence electrons. Two atoms can share an electron, and be held together by their mutual attraction to the shared electron.

18. Covalent Bonds Two fluorine atoms can share each others unpaired valence electron. This results in each atom being attracted to the other’s electron. See figure 23.12 on page 397.

19. Covalent Bonds This type of electrical attraction where atoms are held together by their mutual attraction for shared electrons is called a covalent bond. A substance that is made of atoms with covalent bonds is called a covalent compound.

20. Covalent Bonds

21. Molecules are held together by covalent bonds, so a substance made of molecules is a covalent compound (There are no ions in a covalent compound). Molecules – any group of atoms held together by covalent bonds.

22. Covalent Bonds Elements in the upper-right hand corner of the periodic table (primarily nonmetallic elements) form covalent bonds, because they’re formed by atoms that tend to gain electrons.

23. Covalent Bonds A straight line is used to represent two electrons that are involved in a covalent bond. See page 399 in your text. A double covalent bond or a double bond exists when two atoms share more than two electrons. See page 400 in your text.

24. Covalent Bonds Molecular oxygen, O 2, is an example of this, with four shared electrons. Carbon Dioxide, CO 2, consists of two double bonds connecting oxygen to a central atom. Molecular nitrogen - triple covalent bonds – six electrons – three from each atom are shared.

25. Covalent Bonds Examples of how all these bonds may appear as the following:

26. Covalent Bonds If the electrons in a covalent bond are shared equally, the resulting bond is called a nonpolar covalent bond.

27. Covalent Bonds When one atom pulls the shared electrons toward itself a little more tightly than the other, the resulting covalent bond is said to be a polar bond.

28. Covalent Bonds In a polar bond, the atom that pulls electrons toward itself gains a slight negative charge (because electrons have a negative charge). Since the other atom partially loses an electron, it gains a slight positive charge.

29. Dipoles are molecules that have a slightly positive charge on one end and a slightly negative charge on the other. For example, the atoms in water form polar bonds because oxygen, which has eight protons in its nucleus, has a greater pull on electrons than hydrogen, which has only one proton.