1. Chapter 4 Aqueous Reactions And Solution Stoichiometry

2. Water is soooo abundant! Extremely unique properties High boiling point and melting point Expands upon freezing Essential to support life on earth! Ability to dissolve many substances to form aqueous solutions (only aqueous when water is the dissolving medium)

3. 3 Major types of reactions in aqueous solutions Precipitation reactions Acid-Base reactions Redox reactions

4. But first, How do dissolved substances act in water?

5. Terminology Solution- homogeneous mixture. Solute- what gets dissolved. Solvent- what does the dissolving. Soluble- Can be dissolved. Insoluble- Cannot be dissolved Miscible- liquids dissolve in each other. Immiscible- liquids that do not dissolve in each other

6. Aqueous solutions Dissolved in water. Water is a good solvent because the molecules are polar. The oxygen atoms have a partial negative charge. The hydrogen atoms have a partial positive charge. The angle is 105º

7. Hydration The process of breaking the ions of salts apart. Ions have charges and attract the opposite charges on the water molecules.

8. Hydration H H O H H O H H O H H O H H O H H O H H O H H O H H O

9. Solubility How much of a substance will dissolve in a given amount of water. Usually g/100 mL Varies greatly, but if they do dissolve the ions are separated, and they can move around. Water can also dissolve non-ionic compounds if they have polar bonds.

10. Electrolytes Electricity is moving charges. Electrolytic solutions have the ability to conduct electricity. The ions that are dissolved can move. Solutions of ionic compounds can conduct electricity. (called electrolytic solution) Ionic solids dissociate into it’s component ions as it dissolves

11. 3 Types of solutions Strong electrolytes- completely dissociate (fall apart into many ions). Good conductivity (soluble ionic compounds) Weak electrolytes- Partially fall apart into ions. Due to a smaller number of ions, only slightly conductive Non-electrolytes- Does not break into ions,does not conduct.

12. Molecular Compounds in Water The molecule does not break apart, but remains intact A solution of molecular compounds consists of individual whole molecules dispersed throughout the solutions Since these molecular compounds do not break into ions, they are called non- electrolytes because they do not conduct electricity.

13. Types of solutions Acids- form H + ions when dissolved. Strong acids fall apart completely. many ions H 2 SO 4 HNO 3 HCl HBr HI HClO 4 Weak acids- don’t dissociate completely. Bases - form OH - ions when dissolved. Strong bases- many ions. KOH NaOH

14. Precipitation Reactions When aqueous solutions of ionic compounds are poured together and result in the formation of an insoluble solid (called a precipitate). If you’re not a part of the solution, your part of the precipitate

15. Precipitation reactions NaOH(aq) + FeCl 3 (aq)   NaCl(aq) + Fe(OH) 3 (s) is really Na + (aq)+OH - (aq) + Fe +3 + Cl - (aq)   Na + (aq) + Cl - (aq) + Fe(OH) 3 (s) So all that really happens is OH - (aq) + Fe +3  Fe(OH) 3 (s) Double replacement reaction Since Hydroxide ions are insoluble, the product is a precipitate!

16. Precipitation reaction We can predict the products Can only be certain by experimenting The anion and cation switch partners AgNO 3 (aq) + KCl(aq)  Zn(NO 3 ) 2 (aq) + BaCr 2 O 7 (aq)  CdCl 2 (aq) + Na 2 S(aq) 

17. Precipitations Reactions Only happen if one of the products is insoluble Otherwise all the ions stay in solution- nothing has happened. Need to memorize the rules for solubility (Table 4.1)

18. Solubility Guidelines Solubility- the amount of substance that can be dissolved in a given quantity of solvent Insoluble compounds are those where the attractions between the ions are so large that the water molecule will not separate the ion, and the substance remains mostly undissolved.

19. You can predict it! Since you memorized the solubility guidelines, you can predict whether a substance will produce a precipitate or not.

20. Solubility Rules  All nitrates are soluble  Alkali metals ions and NH 4 + ions are soluble  Halides are soluble except Ag +, Pb +2, and Hg 2 +2  Most sulfates are soluble, except Pb +2, Ba +2, Hg +2,and Sr +2

21. Solubility Rules  Most hydroxides are slightly soluble (insoluble) except compounds of the alkali metal cations, Ca +2, Sr +2, Ba +2  Sulfides, carbonates, chromates, and phosphates are insoluble (see exceptions)  Lower number rules supersede so Na 2 S is soluble

22. Precipitation Reactions Also called double replacement reactions, exchange reactions, or metathesis reactions. The cations exchange anions AgNO 3 + KCl  AgCl + KNO 3

23. Three Types of Equations Molecular Equation- written as whole formulas, not the ions. K 2 CrO 4 (aq) + Ba(NO 3 ) 2 (aq)  Complete Ionic equation show dissolved electrolytes as the ions. 2K + + CrO 4 -2 + Ba +2 + 2 NO 3 -  BaCrO 4 (s) + 2K + + 2 NO 3 - Spectator ions are those that don’t react.

24. Three Type of Equations Net Ionic equations show only those ions that react, not the spectator ions Ba +2 + CrO 4 -2  BaCrO 4 (s) Write the three types of equations for the reactions when these solutions are mixed. iron (III) sulfate and potassium sulfide Lead (II) nitrate and sulfuric acid.

25. Iron III Sulfate + Potassium Sulfide Fe 2 (SO 4 ) 3 + 3K 2 S  3K 2 SO 4 + Fe 2 S 3 2Fe +3 + 3SO 4 -2 + 6K + + 3S -2  6K + + 3SO 4 -2 + Fe 2 S 3 2Fe +3 + 3S -2  Fe 2 S 3

26. Lead II nitrate + Sulfuric acid Pb(NO 3 ) 2 + H 2 SO 4  PbSO 4 + 2HNO 3 Pb +2 + 2NO 3 - + 2H + + SO 4 -2  PbSO 4 + 2H + + NO 3 - Pb +2 + SO 4 -2  PbSO 4

27. Acids Proton donor Monoprotic, diprotic, polyprotic, depending on how many hydrogens it donates. Ionization reactions- acid donates a proton to water to form a hydronium ion. You can only donate 1 proton at a time. HCl (aq)  H + + Cl -

28. Bases Bases are proton acceptors Increases the concentration of OH - ions when added to water Most bases end in OH NH 3 is a base (weak electrolyte

29. Strong vs. Weak Strong acids + strong bases = STRONG ELECTROLYTES (completely ionize in solution) Weak acids + weak bases = WEAK ELECTROLYTES (only partly ionize) Memorize table 4.2 for the complete list of strong acids and bases

30. Acids vs. Bases Acid Base Sour bitter Turn litmus red Turn litmus blue pH 1-6.9 pH 7.1-14

31. Acid-Base Reactions We can consider an acid is a proton donor. a base is a proton acceptor usually OH - What is the net ionic equation for the reaction of HCl(aq) and KOH(aq)? Acid + Base  salt + water H + + OH -  H 2 O

32. Neutralization Reactions An acid and base react to form salt and water Called metathesis reactions because the anion and cation exchange partners

33. Acid – Base Reactions that produce gas The sulfide and carbonate ion act as bases when combine with an acid (they accept a Hydrogen ion) the product of these reactions is a gas HCl + Na 2 S  H 2 S (g) + NaCl HCl + NaHCO 3  NaCl + H 2 CO 3 H 2 CO 3  H 2 O + CO 2 (Used to clean up acid spills) (also used as antacids)

34. Redox Reactions Oxidation-Reduction called Redox An Oxidation-reduction reaction involves the transfer of electrons. In a REDOX Reaction, one substance loses electrons as it is oxidized, the other substance gains electrons and is reduced. Use oxidation numbers to help keep track of where the electrons are in a reaction

35. Oxidation States  The oxidation state of elements in their standard states is zero.  Oxidation state for monoatomic ions are the same as their charge. 3 Oxygen is assigned an oxidation state of -2 in its covalent compounds except as a peroxide.

36. Oxidation states  In compounds with nonmetals hydrogen is assigned the oxidation state +1.  In its compounds fluorine is always –1.  The sum of the oxidation states must be zero in compounds or equal the charge of the ion.

37. Oxidation States Assign the oxidation states to each element in the following. CO 2 NO 3 - H 2 SO 4 Fe 2 O 3 H2OH2O

38. Oxidation-Reduction Transfer electrons, so the oxidation states change. 2Na + Cl 2  2NaCl CH 4 + 2O 2  CO 2 + 2H 2 O Oxidation is the loss of electrons. Reduction is the gain of electrons. OIL RIG LEO GER

39. Oxidation of Metals The reaction of a metal with an acid or a metal salt is a redox reaction (called displacement) Zn + HCl  ZnCl 2 + H 2 Mn + Pb(NO 3 ) 2  Mn(NO 3 ) 2 + Pb

40. Half-Reactions All redox reactions can be thought of as happening in two halves. One produces electrons - Oxidation half. The other requires electrons - Reduction half. Write the half reactions for the following. Na + Cl 2  Na + + Cl - SO 3 - + H + + MnO 4 -  SO 4 - + H 2 O + Mn +2

41. Oxidation-Reduction Oxidation means an increase in oxidation state - lose electrons. Reduction means a decrease in oxidation state - gain electrons. The substance that is oxidized is called the reducing agent. The substance that is reduced is called the oxidizing agent.

42. Reactivity Series Just because it can be written does mot mean it will always happen We can predict whether these displacement reactions will occur by talking about the activity series Look at Table 4.5 The top metals are the most easily oxidized Any metal can be oxidized by the ions below it

43. Opposite for Reference Table This would work Cu + AgCl  Ag + CuCl But this would not Ag + CuCl  Cu + AgCl

44. Agents Oxidizing agent gets reduced. Gains electrons. More negative oxidation state. Reducing agent gets oxidized. Loses electrons. More positive oxidation state.

45. Identify the Oxidizing agent Reducing agent Substance oxidized Substance reduced in the following reactions Fe (s) + O 2 (g)  Fe 2 O 3 (s) Fe 2 O 3 (s)+ 3 CO(g)  2 Fe(l) + 3 CO 2 (g) SO 3 -2 + H + + MnO 4 -  SO 4 -2 + H 2 O + Mn +2

46. Measuring Solutions Concentration- how much is dissolved. Molarity = Moles of solute Liters of solution abbreviated M 1 M = 1 mol solute / 1 liter solution Calculate the molarity of a solution with 34.6 g of NaCl dissolved in 125 mL of solution.

47. Molarity How many grams of HCl would be required to make 50.0 mL of a 2.7 M solution? What would the concentration be if you used 27g of CaCl 2 to make 500. mL of solution?

48. Expressing the Concentration of an Electrolyte 1.0 M solution of NaCl is 1.0 M of Na + ions, and 1.0 M of Cl - ions 1.0 M solution of Na 2 SO 4 is 2M solution of NA + ions, and 1M solution of SO 4 -2 ions 2.0 M solution of Na 2 SO 4 is 4.0 M of Na + ions, and 2.0 M of SO 4 -2 ions

49. Molarity Calculate the concentration of a solution made by dissolving 45.6 g of Fe 2 (SO 4 ) 3 to 475 mL. What is the concentration of each ion?

50. Making solutions Describe how to make 100.0 mL of a 1.0 M K 2 Cr 2 O 4 solution. Describe how to make 250. mL of an 2.0 M copper (II) sulfate dihydrate solution.

51. Dilution Adding more solvent to a known solution. The moles of solute stay the same. moles = M x L M 1 V 1 = M 2 V 2 moles = moles Stock solution is a solution of known concentration used to make more dilute solutions

52. Dilution Adding more solvent to a known solution. The moles of solute stay the same. moles = M x L M 1 V 1 = M 2 V 2 moles = moles Stock solution is a solution of known concentration used to make more dilute solutions

53. Dilution What volume of a 1.7 M solutions is needed to make 250 mL of a 0.50 M solution? 18.5 mL of 2.3 M HCl is added to 250 mL of water. What is the concentration of the solution? 18.5 mL of 2.3 M HCl is diluted to 250 mL with water. What is the concentration of the solution?

54. Dilution You have a 4.0 M stock solution. Describe how to make 1.0L of a.75 M solution. 25 mL 0.67 M of H 2 SO 4 is added to 35 mL of 0.40 M CaCl 2. What mass CaSO 4 Is formed?

55. Acid - Base Reactions Often called a neutralization reaction Because the acid neutralizes the base. Often titrate to determine concentrations. Solution of known concentration (titrant), is added to the unknown (analyte), until the equivalence point is reached where enough titrant has been added to neutralize it.

56. Titration Where the indicator changes color is the endpoint. Not always at the equivalence point. A 50.00 mL sample of aqueous Ca(OH) 2 requires 34.66 mL of 0.0980 M Nitric acid for neutralization. What is [Ca(OH) 2 ]? # of H + x M A x V A = # of OH - x M B x V B