Types of Chemical Reactions and Solution Stoichiometry

1. Types of Chemical Reactionsand Solution Stoichiometry

2. Solutions are homogeneous mixtures Classification of Matter

3. Solute A solute is the dissolved substance in a solution. Salt in salt water Sugar in soda drinks Carbon dioxide in soda drinks Solvent A solvent is the dissolving medium in a solution. Water in salt water Water in soda

4. Saturation of Solutions • A solution that contains the maximum amount of solute that may be dissolved under existing conditions is saturated. • A solution that contains less solute than a saturated solution under existing conditions is unsaturated. • A solution that contains more dissolved solute than a saturated solution under the same conditions is supersaturated.

5. Electrolytes vs. Nonelectrolytes The ammeter measures the flow of electrons (current) through the circuit. • If the ammeter measures a current, and the bulb • glows, then the solution conducts. • If the ammeter fails to measure a current, and the • bulb does not glow, the solution is non-conducting.

6. Definition of Electrolytes and Nonelectrolytes An electrolyte is: • A substance whose aqueous solution conducts • an electric current. A nonelectrolyte is: • A substance whose aqueous solution does not • conduct an electric current. Try to classify the following substances as electrolytes or nonelectrolytes…

7. Electrolytes? • Pure water • Tap water • Sugar solution • Sodium chloride solution • Hydrochloric acid solution • Lactic acid solution • Ethyl alcohol solution • Pure, solid sodium chloride

8. Answers… ELECTROLYTES: NONELECTROLYTES: • Tap water (weak) • NaCl solution • HCl solution • Lactate solution (weak) • Pure water • Sugar solution • Ethanol solution • Pure, solid NaCl But why do some compounds conduct electricity in solution while others do not…?

9. Ionic CompoundsDissociate NaCl(s)  Na+(aq) + Cl-(aq) AgNO3(s)  Ag+(aq) + NO3-(aq) MgCl2(s)  Mg2+(aq) + 2 Cl-(aq) Na2SO4(s)  2 Na+(aq) + SO42-(aq) AlCl3(s)  Al3+(aq) + 3 Cl-(aq)

10. Ions tend to stay in solution where they canconduct a current rather than re-forming a solid. The reason for this is the polar nature of the water molecule… Positive ions associate with the negative end of the water dipole (oxygen). Negative ions associate with the positive end of the water dipole (hydrogen).

11. Some covalent compounds IONIZE in solution Covalent acids form ions in solution, with the help of the water molecules. For instance, hydrogen chloride molecules, which are polar, give up their hydrogens to water, forming chloride ions (Cl-) and hydronium ions (H3O+).

12. Strong acids such as HCl are completelyionized in solution. Other examples of strong acids include: • Sulfuric acid, H2SO4 • Nitric acid, HNO3 • Hydriodic acid, HI • Perchloric acid, HClO4

13. Weak acids such as lactic acid usually ionize less than 5% of the time. Many of these weaker acids are “organic” acids that contain a “carboxyl” group. The carboxyl group does not easily give up its hydrogen.

14. Because of the carboxyl group, organic acids aresometimes called “carboxylic acids”. Other organic acids and their sources include: • Citric acid – citrus fruit • Malic acid – apples • Butyric acid – rancid butter • Amino acids – protein • Nucleic acids – DNA and RNA • Ascorbic acid – Vitamin C This is an enormous group of compounds; these are only a few examples.

15. However, most covalent compounds do not ionizeat all in solution. Sugar (sucrose – C12H22O11), and ethanol (ethyl alcohol – C2H5OH) do not ionize - That is why they are nonelectrolytes!

16. Molarity The concentration of a solution measured in moles of solute per liter of solution. mol= M L

17. Preparation of Molar Solutions Problem: How many grams of sodium chloride are needed to prepare 1.50 liters of 0.500 M NaCl solution? • Step #1: Ask “How Much?” (What volume to prepare?) • Step #2: Ask “How Strong?” (What molarity?) • Step #3: Ask “What does it weigh?” (Molar mass is?) 1.500 L 0.500 mol 58.44 g = 43.8 g 1 L 1 mol

18. Serial Dilution It is not practical to keep solutions of many different concentrations on hand, so chemists prepare more dilute solutions from a more concentrated “stock” solution. Problem: What volume of stock (11.6 M) hydrochloric acid is needed to prepare 250. mL of 3.0 M HCl solution? MstockVstock = MdiluteVdilute (11.6 M)(x Liters) = (3.0 M)(0.250 Liters) x Liters = (3.0 M)(0.250 Liters) 11.6 M = 0.065 L

19. Single Replacement Reactions A + BX  AX + B BX + Y  BY + X Replacement of: • Metals by another metal • Hydrogen in water by a metal • Hydrogen in an acid by a metal • Halogens by more active halogens

20. The Activity Series of the Metals • Lithium • Potassium • Calcium • Sodium • Magnesium • Aluminum • Zinc • Chromium • Iron • Nickel • Lead • Hydrogen • Bismuth • Copper • Mercury • Silver • Platinum • Gold Metals can replace other metals provided that they are above the metal that they are trying to replace. Metals above hydrogen can replace hydrogen in acids. Metals from sodium upward can replace hydrogen in water

21. The Activity Series of the Halogens • Fluorine • Chlorine • Bromine • Iodine Halogens can replace other halogens in compounds, provided that they are above the halogen that they are trying to replace. 2NaCl(s) + F2(g)  2NaF(s) + Cl2(g) ??? MgCl2(s) + Br2(g)  No Reaction ???

22. Double Replacement Reactions The ions of two compounds exchange places in an aqueous solution to form two new compounds. AX + BY  AY + BX One of the compounds formed is usually a precipitate (an insoluble solid), an insoluble gas that bubbles out of solution, or a molecular compound, usually water.

23. Double replacement forming a precipitate… Double replacement (ionic) equation Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq) Complete ionic equation shows compounds as aqueous ions Pb2+(aq) + 2 NO3-(aq) + 2 K+(aq) +2 I-(aq)  PbI2(s) + 2K+(aq) + 2 NO3-(aq) Net ionic equation eliminates the spectator ions Pb2+(aq) + 2 I-(aq)  PbI2(s)

24. Solubility Rules – Mostly Soluble

25. Solubility Rules – Mostly Insoluble

26. Oxidation and Reduction (Redox) • Electrons are transferred • Spontaneous redox rxns can transfer energy • Electrons (electricity) • Heat • Non-spontaneous redox rxns can be made to happen with electricity

27. Oxidation and Reduction An old memory device for oxidation and reduction goes like this… LEO says GER LoseElectrons = Oxidation GainElectrons = Reduction

28. Oxidation Reduction Reactions(Redox) Each sodium atom loses one electron: Each chlorine atom gains one electron:

29. LEO says GER : Lose Electrons = Oxidation Sodium is oxidized Gain Electrons = Reduction Chlorine is reduced

30. Rules for Assigning Oxidation NumbersRules 1 & 2 • The oxidation number of any uncombined element is zero 2. The oxidation number of a monatomic ion equals its charge

31. Rules for Assigning Oxidation NumbersRules 3 & 4 3.The oxidation number of oxygen in compounds is -2 4. The oxidation number of hydrogen in compounds is +1

32. Rules for Assigning Oxidation Number Rule 5 5. The sum of the oxidation numbers in the formula of a compound is 0 2(+1) + (-2) = 0 H O (+2) + 2(-2) + 2(+1) = 0 Ca O H

33. Rules for Assigning Oxidation NumbersRule 6 6. The sum of the oxidation numbers in the formula of a polyatomic ion is equal to its charge X + 4(-2) = -2 S O X + 3(-2) = -1 N O  X = +5  X = +6

34. Reducing Agents and Oxidizing Agents • The substance reduced is the oxidizingagent • The substance oxidized is the reducingagent Sodium is oxidized – it is the reducing agent Chlorine is reduced – it is the oxidizing agent

35. Trends in Oxidation and Reduction • Active metals: • Lose electrons easily • Are easily oxidized • Are strong reducing agents • Active nonmetals: • Gain electrons easily • Are easily reduced • Are strong oxidizing agents

36. Redox Reaction Prediction #1

37. Redox Reaction Prediction #2

38. Not All Reactions are Redox Reactions Reactions in which there has been no change in oxidation number are not redox rxns. Examples: