1. Solution Stoichiometry The concentration of a solution is the amount of solute present in a given quantity of solvent or solution. M = molarity = moles of solute liters of solution What mass of KI is required to make 500. mL of a 2.80 M KI solution? volume KImoles KIgrams KI M KI 500. mL= 232 g KI 166 g KI 1 mol KI x 2.80 mol KI 1 L soln x 1 L 1000 mL x

3. Dilution is the procedure for preparing a less concentrated solution from a more concentrated solution. Dilution Add Solvent Moles of solute before dilution (i) Moles of solute after dilution (f) = MiViMiVi MfVfMfVf =

4. How would you prepare 60.0 mL of 0.200 M HNO 3 from a stock solution of 4.00 M HNO 3 ? M i V i = M f V f M i = 4.00 M f = 0.200V f = 0.0600 L V i = ? L V i = MfVfMfVf MiMi = 0.200 x 0.0600 4.00 = 0.00300 L = 3.00 mL 3.00 mL of acid + 57.0 mL of water= 60.0 mL of solution

5. Titrations In a titration a solution of accurately known concentration is gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete. Equivalence point – the point at which the reaction is complete Indicator – substance that changes color at (or near) the equivalence point Slowly add base to unknown acid UNTIL the indicator changes color

6. What volume of a 1.420 M NaOH solution is Required to titrate 25.00 mL of a 4.50 M H 2 SO 4 solution? WRITE THE CHEMICAL EQUATION! volume acidmoles acidmoles basevolume base H 2 SO 4 + 2NaOH 2H 2 O + Na 2 SO 4 4.50 mol H 2 SO 4 1000 mL soln x 2 mol NaOH 1 mol H 2 SO 4 x 1000 ml soln 1.420 mol NaOH x 25.00 mL = 158 mL M acid rx coef. M base

7. What volume of a 0.800 M HCl solution is required to titrate 40.00 mL of a 1.600 M NH 3 solution? HCl + NH 3 NH 4 Cl 1.600 mol NH 3 1000 mL soln x 1 mol HCl 1 mol NH 3 x 1000 ml soln 0.800 mol HCl x 40.00 mL = 80.0 mL

8. What volume of a 1.200 M Ba(OH) 2 solution is required to titrate 90.00 mL of a 3.600 M H 3 PO 4 solution? 3Ba(OH) 2 + 2H 3 PO 4 6H 2 O + Ba 3 (PO 4 ) 2 3.600 mol H 3 PO 4 1000 mL soln x 3 mol Ba 2 (OH) 2 mol H 3 PO 4 x 1000 ml soln 1.200 mol Ba 2 (OH) x 90.00 mL = 405.0 mL

9. Reactions in Aqueous Solution Chapter 4 71% of Earth’s surface is water Another 3% is covered with ice 66% of human body is water

10. A solution is a homogenous mixture of 2 or more substances The solute is(are) the substance(s) present in the smaller amount(s) The solvent is the substance present in the larger amount SolutionSolventSolute Soft drink (l) Air (g) Soft Solder (s) H2OH2O N2N2 Pb Sugar, CO 2 O 2, Ar, CH 4 Sn

11. An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity. A nonelectrolyte is a substance that, when dissolved, results in a solution that does not conduct electricity. nonelectrolyte weak electrolyte strong electrolyte

12. Strong Electrolyte – 100% dissociation NaCl (s) Na + (aq) + Cl - (aq) H2OH2O Weak Electrolyte – not completely dissociated CH 3 COOH CH 3 COO - (aq) + H + (aq) Electrolyte Conducts electricity in solution? Cations (+) and Anions (-) Nonelectrolyte does not conduct electricity? No cations (+) and anions (-) in solution C 6 H 12 O 6 (s) C 6 H 12 O 6 (aq) H2OH2O

14. Table 1 Electrolyte Classification of Some Common Substances Strong ElectrolytesWeak ElectrolytesNonelectrolytes HCl, HBr, HICH 3 COOHH 2 O HClO 4 HFCH 3 OH HNO 3 C 2 H 5 OH H 2 SO 4 C 12 H 22 O 11 (sucrose) KBrMost organic compounds NaCl NaOH, KOH Other soluble ionic compounds E.g. Calculating ion concentrations in a solution of strong electrolyte. What are the concentration (in molarity) of Mg 2+ and Cl - ions in 5.0g/L of MgCl 2 solution? [Mg 2+ ] = 5.0g x 1 molMgCl 2 x 1 mol Mg 2+ =5.3 x10 -2 mol/L L 95.21gMgCl 2 1 molMgCl 2 [Cl - ] =2 x 5.3x10 -2 = 0.11 mol/L

15. Concentration of Ions A bottle labeled as 0.100 M Al 2 (SO 4 ) 3. [Al 3+ ] = _____ M (mol / L) [SO 4 2– ] = _____ M Assume sea water is 0.438 M NaCl, 0.0512 M MgCl 2, and 0.001 M CaCl 2 [Na + ] = _____ M [Mg 2+ ] = _____ M [Ca 2+ ] = _____ M [Cl – ] = _____ M Know how to calculate your quantities

16. Aqueous Reactions Aqueous reactions can be grouped into three general categories, each with its own kind of driving force: 1. Precipitation reactions 2. Acid base neutralization reactions 3. Oxidation-reduction reactions.

17. Precipitation Reactions When ions form a solid that is not very soluble, a solid is formed. Such a phenomenon is called precipitation. Precipitate – insoluble solid that separates from solution PbI 2

18. Rules for converting molecular equations to ionic equations The rules for converting molecular equations to ionic equations follow: 1) Make sure the molecular equation is balanced 2) Ionic substances indicated in the molecular equation as dissolved in solution, such as NaCl(aq), are normally written as ions. 3) Ionic substances that are insoluble (do not dissolve) either as reactants or products (such as precipitate) are represented by formulas of the compounds 4) Molecular substances that are strong electrolytes, such as strong acids, are written as ions. Thus, HCl(aq) is written as H 3 O + (aq) + Cl - (aq) or as H + (aq) + Cl - (aq). 5) Molecular substances that are weak electrolytes or nonelectrolytes are represented by their molecular formulas.

19. 1. All alkali metals (Group 1A) compounds are soluble. Solubility Rules for Common Ionic Compounds In water at 25 o C 2. All ammonium (NH 4 + ) compounds are soluble. 3. All nitrate (NO 3 - ), chlorate (ClO 3 - ), and perchlorate (ClO 4 - ) compounds are soluble. 4. Most hydroxides (OH - ) are insoluble. The exceptions are barium hydroxide [Ba(OH) 2 ], which is very soluble, calcium hydroxide [Ca(OH) 2 ], which is slightly soluble, and previous examples. 5. Most compounds containing chlorides (Cl - ), bromides (Br - ), or iodides (I - ), are soluble. The exceptions are those containing Ag +, Hg 2 2+, and Pb 2+. 7. All carbonates (CO 3 2- ), phosphates (PO 4 3- ), and sulfides (S 2- ) are insoluble, excepting previous rules. 6. Most sulfates (SO 4 2- ) are soluble, excepting previous rules. Calcium sulfate [CaSO 4 ] and silver sulfate [Ag 2 SO 4 ] are slightly soluble. Barium sulfate [BaSO 4 ], mercury (II) sulfate [HgSO 4 ], and lead sulfate [PbSO 4 ] are insoluble.

20. Writing Net Ionic Equations 1.Write the balanced molecular equation. 2.Write the ionic equation showing the strong electrolytes 3.Determine precipitate from solubility rules 4.Cancel the spectator ions on both sides of the ionic equation AgNO 3 (aq) + NaCl (aq) AgCl (s) + NaNO 3 (aq) Ag + + NO 3 - + Na + + Cl - AgCl (s) + Na + + NO 3 - Ag + + Cl - AgCl (s) Write the net ionic equation for the reaction of silver nitrate with sodium chloride.

21. write a net ionic equation for the reaction. (a) Al 2 (SO 4 ) 3 + NaOH  i) write down the reactants and interchange of anions to get product Al 2 (SO 4 ) 3 + 6NaOH  2Al(OH) 3 + 3Na 2 SO 4 All common Na compounds are water soluble Na + remain in solution. The combination of Al 3+ and OH - produce insoluble Al(OH) 3. Then the ionic equation is 2Al 3+ +3SO 4 2- + 6Na + + 6OH -  2Al(OH) 3 (s)+ 6Na + + 3SO 4 2- The net ionic equation is : Al 3+ + 3OH -  Al(OH) 3 (s)

22. Net Ionic Equations Note that some ions appear on both side of equation. These ions go through the reaction unchanged- does not take part in the reaction. We called them spectator ions. We can cancel them from the equation. The resulting equation is a net ionic equation. Ag + (aq) + I-(aq)  AgI(s) Net ionic Equation A net ionic equation is an equation that includes only the actual participants in a reaction, with each participant denoted by the symbol or formula that best represent it.

23. Precipitation Reactions Precipitate – insoluble solid that separates from solution molecular equation ionic equation net ionic equation Pb 2+ + 2NO 3 - + 2Na + + 2I - PbI 2 (s) + 2Na + + 2NO 3 - Na + and NO 3 - are spectator ions PbI 2 Pb(NO 3 ) 2 (aq) + 2NaI (aq) PbI 2 (s) + 2NaNO 3 (aq) precipitate Pb 2+ + 2I - PbI 2 (s)

24. Precipitation Reactions Ag + (aq) + NO 3 – (aq) + Cs + (aq) + I – (aq)  AgI (s) + NO 3 – (aq) + Cs + (aq) Ag + (aq) + I – (aq)  AgI (s) (net reaction) or Ag + + I –  AgI (s) Heterogeneous Reactions Spectator ions Soluble ions Alkali metals, NH 4 + nitrates, ClO 4 -, acetate Mostly soluble ions Halides, sulfates Mostly insoluble Silver halides Metal sulfides, hydroxides carbonates, phosphates

25. Acids (according to Arrhenius) Have a sour taste. Vinegar owes its taste to acetic acid. Citrus fruits contain citric acid. React with certain metals to produce hydrogen gas. React with carbonates (CO 3 2- ) and bicarbonates (HCO 3 - ) to produce carbon dioxide gas. Are electrolytes. Cause certain plant dyes to change color. For example, they turn litmus from blue to red.

26. Have a bitter taste. Feel slippery (they make water insoluble organic molecules into water soluble molecules). Many soaps contain bases. Bases (according to Arrhenius) Are electrolytes. Cause certain plant dyes to change color. For example, they turn litmus from red to blue.

27. Arrhenius acid is a substance that produces H + (H 3 O + ) in water Arrhenius base is a substance that produces OH - in water

28. A Brønsted acid is a proton donor A Brønsted base is a proton acceptor acidbaseacidbase A Brønsted acid must contain at least one ionizable proton!

29. Monoprotic acids HCl H + + Cl - HNO 3 H + + NO 3 - CH 3 COOH H + + CH 3 COO - Strong electrolyte, strong acid Weak electrolyte, weak acid Diprotic acids H 2 SO 4 H + + HSO 4 - HSO 4 - H + + SO 4 2- Strong electrolyte, strong acid Weak electrolyte, weak acid Triprotic acids H 3 PO 4 H + + H 2 PO 4 - H 2 PO 4 - H + + HPO 4 2- HPO 4 2- H + + PO 4 3- Weak electrolyte, weak acid

30. 2. Neutralization Reaction acid + base salt + water HCl (aq) + NaOH (aq) NaCl (aq) + H 2 O H + + Cl - + Na + + OH - Na + + Cl - + H 2 O H + + OH - H 2 O net ionic equation for strong acid / strong base reaction

31. Neutralization Reaction acid + base salt + water HCl (aq) + NH 3 (aq) NH 4 Cl(aq) H + + Cl - + NH 3 (aq) NH 4 + + Cl - H + + NH 3 (aq) NH 4 + net ionic equation for strong acid / weak base reaction

32. Oxidation is the lose of one or more electrons by a substance. Historically oxidation is combination of an element with oxygen Reduction is the gain of one or more electrons by a substance. R eduction is referred to the removal of oxygen from an oxide A redox reaction is a process in which electrons are transferred between substance or in which atoms change oxidation number. Oxidizing agent: contains an element whose oxidation state decreases in a redox reaction (it make possible for some other substance to be oxidized and itself reduced.) gains electrons (electrons are found on the left side of its half-equation). Reducing agent: contains an element whose oxidation state increases in a redox reaction, loses electrons (electrons are found on the left side of its half-equation). Oxidation-Reduction Reaction

33. Oxidation number The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred. 1.Free elements (uncombined state) have an oxidation number of zero. Na, Be, K, Pb, H 2, O 2, P 4 = 0 2.In monatomic ions, the oxidation number is equal to the charge on the ion. Li +, Li = +1; Fe 3+, Fe = +3; O 2-, O = -2 3.The oxidation number of oxygen is usually –2. In H 2 O 2 and O 2 2- it is –1.

34. 4.The oxidation number of hydrogen is +1 except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1. 6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion. 5.Group IA metals are +1, IIA metals are +2 and fluorine is always –1. HCO 3 - O = -2H = +1 3x(-2) + 1 + ? = -1 C = +4 Oxidation numbers of all the elements in HCO 3 - ?

35. Fig. 4.10

36. NaIO 3 Na = +1 O = -2 3x(-2) + 1 + ? = 0 I = +5 IF 3 F = -1 3x(-1) + ? = 0 I = +3 K 2 Cr 2 O 7 O = -2K = +1 7x(-2) + 2x(+1) + 2x(?) = 0 Cr = +6 Oxidation numbers of all the elements in the following ?

37. Oxidation-Reduction Reactions (electron transfer reactions) 2Mg (s) + O 2 (g) 2MgO (s) 2Mg 2Mg 2+ + 4e - O 2 + 4e - 2O 2- Oxidation half-reaction (lose e - ) Reduction half-reaction (gain e - ) 2Mg + O 2 + 4e - 2Mg 2+ + 2O 2- + 4e - 2Mg + O 2 2MgO

38. Fe (s) + 2HCl (aq) FeCl 2 (aq) + H 2 (g) Fe is oxidizedFe Fe 2+ + 2e - H + is reduced2H + + 2e - H 2 Fe is the reducing agent H + is the oxidizing agent Iron (Fe) reacting with hydrochloric acid (HCl) An oxidizing agent: contains an element whose oxidation state decreases in a redox reaction (it make possible for some other substance to be oxidized and itself reduced.) gains electrons (electrons are found on the left side of its half-equation). A reducing agent: contains an element whose oxidation state increases in a redox reation loses electrons (electrons are found on the left side of its half-equation).

40. Zn (s) + CuSO 4 (aq) ZnSO 4 (aq) + Cu (s) Zn is oxidizedZn Zn 2+ + 2e - Cu 2+ is reducedCu 2+ + 2e - Cu Zn is the reducing agent Cu 2+ is the oxidizing agent Copper wire reacts with silver nitrate to form silver metal. What is the oxidizing agent in the reaction? Cu (s) + 2AgNO 3 (aq) Cu(NO 3 ) 2 (aq) + 2Ag (s) Cu Cu 2+ + 2e - Ag + + 1e - AgAg + is reducedAg + is the oxidizing agent Cu is oxidizedCu is the reducing agent

41. Ca 2+ + CO 3 2- CaCO 3 NH 3 + H + NH 4 + Zn + 2HCl ZnCl 2 + H 2 Ca + F 2 CaF 2 Precipitation Acid-Base Redox (H 2 Displacement) Redox (Combination) Classify the following reactions.

42. Chapter 4: Reactions in Aqueous Solutions