The Chemical Bond

1. The Chemical Bond Ionic and Covalent Bonds Mr. Nelson Chemistry

2. Ionic Compounds Overview • Bonds formed between a metal and a nonmetal • Metals form positive ions called cations • Trick to remember: “ca+ion” • Non-metals form negative ions called anions • Ionic bonds are formed by a transfer of electrons

3. Properties of Ionic Compounds • Ionic compounds dissociate (or break apart) in liquids • Example: Table salt in water • Ionic compounds conduct electricity • If something conducts electricity, then ions/charges must be able to move • Solid ionic compounds are rigid, not allowing charges to move • Two conditions allow ionic compounds to conduct electricity: • Molten (or liquid) • Dissolved in solution • Example: Li-ion batteries

4. Properties of Ionic Compounds • Ionic compounds tend to have high melting points • This is due to the high strength of the ionic bonds • Stronger bonds need more energy to break apart!

5. Formation of Ionic Compounds • Ionic bonds follow the Octet Rule • Octet rule – Each element tends to form compounds based on the easiest way to gain 8 valence electrons • Example: Na and Cl • Sometimes combinations are not an easy one-to-one ratio! • Example: Ca and Cl

6. Formation of Ionic Compounds • Examples • Mg and O • K and S • Practice • Ba and N • Li and P

7. Formation of Ionic Compounds • Transition metal ionic compounds • All form positive ions (cations), but can form multiple ions • Example: FeCl2 and FeCl3

8. Covalent Compounds Overview • Covalent bonds are made between nonmetals only • Does NOT depend on charges!! • Covalent bonds will share electrons as both elements need to gain electrons to obtain full outer shell (8 valence electrons)

9. Properties of Covalent Compounds • Forms longer bonds because bonds are weaker than ionic bonds • Energy is lowered when atoms form a covalent bond • This is the main driving force behind their formation • There is a balance between their attractive and repulsive forces • Potential energy curve for H2 bond

10. Potential Energy Curve of H2 Bond Potential Energy Bond Length

11. Two Types of Covalent Bonds • Polar covalent bonds are covalent bonds where electrons are shared unequally between the atoms • Non-polar covalent bonds are covalent bonds where electrons are shared equally between the atoms

12. Polarity • Polarity depends on differences in electronegativity • The greater the difference in electronegativity, the more polar the bond is • Example: Cl2 and HF • The more polar the bond, the stronger the bond • Polar covalent bonds have stronger bonds than non-polar bonds • Stronger bonds give higher boiling points (just like ionic compounds)

13. The Chemical Bond Overview • In order of increasing bond strength: • In order of increasing bond length:

14. Naming Ionic Compounds • Simple Naming Rules • Identify the compound as Ionic (Metal and a non-metal) • Determine charges for each ion • Overall charge of compound must be zero • General formula for names: • Cation + Anion Root-ide

15. Examples • Calcium fluoride • Calcium phosphide • Magnesium nitride • Potassium oxide

16. Examples • LiCl2 • SrBr2 • BaI2

17. Transition Metal Compounds • Since transition metals form multiple ions, roman numerals denote charge Roman numerals are not the number of molecules! No, really. Not the molecules

18. Examples • iron (III) chloride • copper (II) oxide • iron (III) oxide • silver (I) sulfide

19. Polyatomic Ion Compounds • Some elements are considered as a charged group of two or more atoms. They are considered as a single ion called a polyatomic ion. • You will not need to memorize these, only recognize it when you see them

20. Examples • Aluminum chromate • Potassium hydroxide • Sodium thiosulfate • Calcium acetate