Lab #5 Chemical Bonds & Bonding

1. Lab #5Chemical Bonds & Bonding Chemistry 108 Instructor:Robert Goldman

2. Monatomic ions • Charged atom • AKA monatomic ion • Results from a different number of protons and electrons.

3. Polyatomic ion • A positively or negatively charged compound. • The atoms in these compounds do not “cancel out” each other’s charges. In most compounds, the atoms charges sum to “0”.

4. Polar vs. Non-polar • “Balanced or Unbalanced” or “Symmetrical” • Polar molecules are unbalanced or unsymmetrical. • Non-polar molecules are balanced or symmetrical.

5. Polar vs. Non-polar • What causes “unbalance” in a molecule? • DIFFERENCES IN ELECTRONEGATIVITY! H F F F

6. Polar vs. Non-polar F F H F + -

7. Polar vs. Non-polar H H C H H

8. Solvents • A solvent is a substance in a solution that dissolves the other components • Sometimes also referred to as the major component, or substance present in the greatest amount.

9. Solutes • Component which becomes dissolved in the solvent. • Sometimes referred to as the “minor component”; present in smaller amounts.

10. Precipitate • Products formed in a reaction that are not soluble in the solvent being used. • These appear as small solids floating in the solvent.

11. Polar and Non-polarSolvents/Solutes • Polar solvents will dissolve polar solutes. • Non-polar solvents will dissolve non-polar solutes. • REMEMBER: LIKE DISSOLVES LIKE!

12. Emulsifiers • An emulsifier is a substance that allows non-polar solvents dissolve polar solutes and vice-versa. • Dish soap is an emulsifier. It facilitates the solubility of oily grime to remove it from dishes.

13. Ionic Equations • Emphasize the existence of any water soluble ions in solution and show their presence. • Net ionic equations are derived from these and focus on any ions which are undergoing dramatic changes during the chemical reaction.

14. Standard  Ionic equations Example: NaBr + AgNO3  NaNO3 +AgBr Step 1: Check solubility chart on p.234 of the textbook for solubilities. (Everything here is soluble BUT AgBr.)

15. Standard  Ionic equations Example: Separate the soluble compounds to illustrate their disassociation in water. NaBr + AgNO3  NaNO3 +AgBr becomes Na+ + Br- + Ag+ + NO3-  Na + + NO3- +AgBr

16. Standard  Ionic equations • Some ionic compounds are shown a bit differently: • MgF2 is shown as Mg2+ + 2 F- • K2SO4 is shown as 2 K+ + SO4-2

17. Net Ionic Equations • Remove all spectator ions (those that appear on both sides) from the equation. Example: Mg+2 + 2 NO3- + 2 Na+ + CO3-2 MgCO3 + 2 Na++ 2 NO3- becomes Mg+2 + CO3-2 MgCO3

18. Covalent Bond Formation • Atoms share electrons to attain a lower energy configuration (octet in most cases). • Example: F with 7e- covalently bonds with Br with 7e-, bromine, electrons are shared, but unequally due to different electronegativities. • Pure covalent bonds are formed by identical atoms which share electrons and are nonpolar. • GENERALLY TWO NON-METALLIC ELEMENTS WILL FORM COVALENT BONDS.

19. Ionic Bond Formation • Ionic bonds are formed when an electronegative element, which has a high electron affinity, steals an electron from a metal, which has a low electron affinity and is willing to give it up. Example: Na+ + Cl- form NaCl (overall neutrally charged) (Neighboring atoms of opposite charge attract) Polar bonds are always formed in this case.

20. Today in Lab • Test solutions at conductivity station. Current only flows (and bulb only lights) if ions are present as they are required to conduct electricity. Be sure to rinse the electrodes between tests. • Test solubilites of common salts in polar vs. non-polar solvent.

21. Today in Lab • Using the solutions created in part B, combine as directed, observe results, and complete equations. • STOICHIOMETRY IS NEXT WEEK - BRING A CALCULATOR!