1. Chapter 4: Reactions in Aqueous Solutions Chemistry 1411 Joanna Sabey 1

2. Aqueous Solutions Solution: a homogeneous mixture of two or more substances. Solute: the substance present in the smaller amount. Solvent: the substance present in a larger amount. Aqueous solutions: the solute initially is a liquid or solid and the solvent is water. 2

3. Electrolytes Electrolyte: a substance that, when dissolved in water, results in a solution that can conduct electricity. Nonelectrolyte: does not conduct electricity when dissolved in water. 3

4. Electrolytes Strong Electrolytes will dissociate completely in water. Break down into anions and cations. NaCl  Na + (aq) + Cl - (aq) Hydration: the process in which an ion is surrounded by water molecules arranged in a specific manner. 4

5. Electrolytes Weak Electrolytes: These compounds do not ionize completely and have a reversible reaction. CH 3 COOH CH 3 COO - (aq) + H + (aq) 5 StrongWeakNonelectrolyte HClCH 3 COOH(NH 2 ) 2 CO HNO 3 HFCH 3 OH HClO 4 HNO 2 C 2 H 5 OH H 2 SO 4 NH 3 C 6 H 12 O 6 NaOHC 12 H 22 O 11 Ionic Compounds

6. Precipitation Reactions Precipitate: an insoluble solid that separate from a solution. Precipitation reaction: a reaction that occurs in aqueous solution which the result it the formation of a precipitate. Pb(NO 3 ) 2 (aq) + 2 KI(aq)  PbI 2 (s) + 2 KNO 3 (aq) Metathesis reaction(double displacement): a reaction that involves the exchanging of parts between the two compounds. AX + BY  AY + BX 6

7. Precipitation Reactions Solubility: the maximum amount of solute that will dissolve in a give quantity of solvent at a specific temperature. 7 Rule StatementExceptions 1All group 1A and ammonium compounds are soluble. - 2All nitrates are soluble.- 3Acetates are soluble.- 4Most sulfates are soluble.CaSO 4, SrSO 4,BaSO 4, Ag 2 SO 4, Hg 2 (SO 4 ) 2,PbSO 4 5Most halides are solubleAg +, Hg 2 2+, Pb 2+ 6Most carbonates are insoluble.Group IA, (NH 4 ) 2 CO 3 7Most phosphates are insoluble.Group IA, (NH 4 ) 3 PO 4 8Most sulfides are insoluble.Group IA, (NH 4 ) 2 S 9Most hydroxide are insoluble.Group IA, Sr(OH) 2, Ba(OH) 2 10Most chromates are insoluble.Group IA, (NH 4 ) 2 CrO 4

8. Solubility Sodium sulfate, Na 2 SO 4 ◦ Soluble Aluminum nitrate, Al(NO 3 ) 3 ◦ Soluble Barium sulfate, BaSO 4 ◦ Insoluble Lead (II) chromate,PbCrO 4 ◦ insoluble Ammonium sulfide, (NH 4 ) 2 S ◦ Soluble 8

9. Double-Displacement Equations Molecular Equation: The formulas of the compounds are written as though all species existed as molecules or whole units. If a precipitate does not form, a double- displacement reaction does not take place. The reaction of lead (II) nitrate and potassium iodide. Pb(NO 3 ) 2 (aq) + 2 KI(aq)  PbI 2 (s) + 2 KNO 3 (aq) 9

10. Ionic Equations Ionic Equations: shows dissolved species as free ions. Spectator ions: ions that are not involved in the overall reaction. Ionic Equation: Pb 2+ (aq) + 2 NO 3 - (aq) +2 K + (aq) + 2 I - (aq)  PbI 2 (s) + 2 K + (aq) + 2 NO 3 - (aq) 10

11. Ionic Equations Net Ionic Equation: shows only the species that actually take part in the reaction. Pb 2+ (aq) + 2 I - (aq)  PbI 2 (s) 11

12. Double-Displacement Reactions Write the balanced molecular equation. Write the ionic equation showing the strong electrolytes, aqueous compounds, completely dissociated into cations and anions. Cancel the spectator ions on both sides of the ionic equation. Check that charges and number of atoms are balanced in the net ionic equation. 12

13. Double-Displacement Reactions Write the molecular equation, the ionic equation and the net ionic equation for the reaction of aqueous barium chloride with aqueous potassium chromate. What are we starting with? ◦ BaCl 2 (aq) and K 2 CrO 4 (aq) What are we producing? ◦ KCl (aq) and BaCrO 4 (s) Write the molecular equation: ◦ BaCl 2 (aq) + K 2 CrO 4 (aq)  KCl (aq) + BaCrO 4 (s) Balance the molecular equation: ◦ BaCl 2 (aq) + K 2 CrO 4 (aq)  2 KCl (aq) + BaCrO 4 (s) Break down the aqueous compounds into ions, ionic equation: ◦ Ba 2+ (aq) + 2 Cl - (aq) + 2 K + (aq) + CrO 4 2- (aq)  2 K + (aq) + 2 Cl - (aq) + BaCrO 4 (s) Write the net ionic equation: ◦ Ba 2+ (aq) + CrO 4 2- (aq)  BaCrO 4 (s) 13

14. Double-Displacement Reactions Write the molecular equation, the ionic equation, and the net ionic equation of the reaction of aqueous potassium phosphate with aqueous calcium nitrate. What are we starting with? ◦ K 3 PO 4 (aq) and Ca(NO 3 ) 2 (aq) What are we producing? ◦ KNO 3 (aq) and Ca 3 (PO 4 ) 2 (s) Write the molecular equation: ◦ K 3 PO 4 (aq) + Ca(NO 3 ) 2 (aq)  Ca 3 (PO 4 ) 2 (s) + KNO 3 (aq) Balance the molecular equation: ◦ 2 K 3 PO 4 (aq) + 3 Ca(NO 3 ) 2 (aq)  Ca 3 (PO 4 ) 2 (s) + 6 KNO 3 (aq) Break down the aqueous compounds into ions, ionic equation: ◦ 6 K + (aq) + 2 PO 4 3- (aq) + 3 Ca 2+ (aq) + 6 NO 3 - (aq)  Ca 3 (PO 4 ) 2 (s) + 6 K + (aq) + 6 NO 3 - (aq) Write the net ionic equation: ◦ 2 PO 4 3- (aq) + 3 Ca 2+ (aq)  Ca 3 (PO 4 ) 2 (s) 14

15. Acid-Base Reactions An acid and a base react. Acids: ◦ Have a sour taste. ◦ Cause color changes in plant dyes ◦ React with certain metals to produce hydrogen gas.  2HCl (aq) + Mg (s)  MgCl 2 (aq) + H 2 (g) ◦ React with carbonates and bicarbonates to produce carbon dioxide gas.  2HCl (aq) + CaCO 3 (s)  CaCl 2 (aq) + CO 2 (g) + H 2 O (l) ◦ Aqueous acid solutions conduct electricity. 15

16. Acid-Base Reactions Base: ◦ Have a bitter taste. ◦ Feel slippery. Many soaps contain bases. ◦ Cause color changes in plant dyes. ◦ Aqueous base solutions conduct electricity. 16

17. Acid-Base Reactions Brønsted Acid: proton donor. Brønsted Base: proton acceptor. HCl(aq) + H 2 O(l)  H 3 O + (aq) + Cl - (aq) ◦ Acid: HCl ◦ Base: H 2 O 17

18. Acid-Base Reactions Monoprotic Acids: each unit of the acid yields one hydrogen ion upon ionization. ◦ HCl  H + + Cl - Diprotic Acids: each unit of the acid gives up to hydrogen ions in two separate steps. ◦ H 2 SO 4  H + + HSO 4 - ◦ HSO 4 -  H + + SO 4 2- Triprotic acids: each unit of the acid gives three hydrogen ions in three separate steps. ◦ H 3 PO 4  H + + H 2 PO 4 - ◦ H 2 PO 4 -  H + + HPO 4 2- ◦ HPO 4 2-  H + + PO 4 3- 18

19. Neutralization reactions Neutralization reaction: a reaction between an acid and a base to produce water and a salt. HCl(aq) + NaOH (aq)  NaCl(aq) + H 2 O (l) Write the molecular, ionic and net ionic equation for the acid-base reaction of hydrobromic acid and barium hydroxide. Write the molecular equation: ◦ HBr(aq) + Ba(OH) 2 (aq)  BaBr 2 (aq) + H 2 O(l) Balance the molecular equation: ◦ 2 HBr(aq) + Ba(OH) 2 (aq)  BaBr 2 (aq) + 2 H 2 O(l) Write the ionic equation: ◦ 2 H + (aq) + 2 Br - (aq) + Ba 2+ (aq) + 2 OH - (aq)  Ba 2+ (aq) + 2 Br - (aq) + 2 H 2 O(l) Write the net ionic equation: ◦ 2 H + (aq) + 2 OH - (aq)  2 H 2 O(l) 19

20. Neutralization Reactions Write the molecular, ionic and net ionic equation for the acid-base reaction of sulfuric acid and potassium hydroxide. Write the molecular equation: ◦ H 2 SO 4 (aq) + KOH(aq)  K 2 SO 4 (aq) + H 2 O(l) Balance the molecular equation: ◦ H 2 SO 4 (aq) + 2 KOH(aq)  K 2 SO 4 (aq) + 2 H 2 O(l) Write the ionic equation: ◦ 2 H + (aq) + SO 4 2 - (aq) + 2 K + (aq) + 2 OH - (aq)  2 K + (aq) + SO 4 2 - (aq) + 2 H 2 O(l) Write the net ionic equation: ◦ 2 H + (aq) + 2 OH - (aq)  2 H 2 O(l) 20

21. Oxidation-Reduction Reactions Oxidation-Reduction reactions: redox reactions, electron transfer reactions. 2 Mg(s) + O 2 (g)  2 MgO(s) Oxidation reaction: 2 Mg  2 Mg 2+ + 4e- Reduction reaction: O 2 + 4e-  2 O 2- OIL RIG 21

22. Oxidation-Reduction Reaction Reducing Agent: donates electrons Oxidizing Agent: accepts electrons Oxidation Number: the number of charges the atom would have in a molecule if electrons were transferred completely. 0 0 +1-1 H 2 (g) + Cl 2 (g)  2 HCl(g) 22

23. Oxidation Numbers Free elements (uncombined state) have an oxidation number of zero. In monatomic ions, the oxidation number is equal to the charge on the ion. The oxidation number of oxygen is usually –2. In H 2 O 2 and O 2 2- it is –1. 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. Group IA metals are +1, IIA metals are +2 and fluorine is always –1 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. 23

24. Oxidation Numbers Assign oxidation numbers to all the elements in the following compounds and ions: Li 2 O ◦ Li=+1 X 2 = +2 ◦ O= -2 HNO 3 ◦ H= +1 ◦ O= -2 X 3 = -6 ◦ N is left over and we want the overall charge to be zero: ◦ N= +5 Cr 2 O 7 2- ◦ O = -2 X 7 = -14 ◦ Cr is left over and we want the overall charge to be -2: ◦ Cr = 2 x ? = +12 ◦ Cr = +6 24

25. Redox Reactions Combination reaction: redox reaction in which two or more substances combine to form a single product. A + B  C S(s) + O 2 (g)  SO 2 (g) 0 0 +4-2 Decomposition reaction: redox reaction in which a compound breaks down into two or more components C  A + B 2 HgO(s)  2 Hg(l) + O 2 (g) +2-2 0 0 25

26. Redox Reactions Combustion reaction: redox reaction in which a substance reacts with oxygen usually with the release of heat and light. C 3 H 8 (g) + 5 O 2 (g)  3 CO 2 (g) + 4 H 2 O(l) 0 -2 -2 26

27. Displacement Reactions Displacement Reaction: an ion, or atom, in a compound is replaced by an ion, or atom, of another element. 27

28. Displacement Reactions Hydrogen Displacement: All alkali metals and some alkaline earth metals, will displace hydrogen from cold water. 2 Na(s) + 2 H 2 O(l)  2 NaOH(aq) + H 2 (g) 0 +1-2 +1-2+1 0 Metal Displacement: One metal in a compound can be replaced by another metal in the elemental state. Based on Metal Activity Series FeSO 4 (aq) + Co (s)  No Reaction V 2 O 5 (s) + 5 Ca(l)  2 V(l) + 5 CaO(s) +5-2 0 0 +2-2 28

29. Displacement Reactions Halogen Displacement: A halogen takes the place of another halogen. F 2 > Cl 2 > Br 2 > I 2 Cl 2 (g) + 2 KBr(aq)  2 KCl + Br 2 (l) 0 +1-1 +1-1 0 Cl 2 (g) + KF(aq)  N.R. 29

30. Disproportionation Reaction Disproportionation Reaction: an element in one oxidation state is simultaneously oxidized and reduced. 2 H 2 O 2 (aq)  2 H 2 O(l) + O 2 (g) +1-1 +1-2 0 Cl 2 (g) +2 OH - (aq)  ClO - (aq) + Cl - (aq) + H 2 O(l) 0 -2+1 +1-2 -1 +1-2 30

31. Redox Reactions Classify the following Redox Reactions: 2 N 2 O(g)  2 N 2 (g) + O 2 (g) ◦ Assign oxidation numbers  N: +1  0  O: -2  0 ◦ State type of reaction:  Decomposition reaction 6 Li(s) + N 2 (g)  2 Li 3 N(s) ◦ Assign oxidation numbers:  Li: 0  +1  N: 0  -3 ◦ State type of reaction:  Combination reaction 31

32. Redox Reactions Classify the following reactions: Ni(s) + Pb(NO 3 ) 2 (aq)  Pb(s) + Ni(NO 3 ) 2 (aq) ◦ Assign oxidation numbers:  Ni: 0  +2  Pb: +2  0 ◦ State type of reaction:  Metal displacement reaction 2 NO 2 (g) + H 2 O(l)  HNO 2 (aq) + HNO 3 (aq) ◦ Assign oxidation numbers:  N: +4  +3 and +5 H: +1  +1  O: -2  -2 ◦ State type of reaction:  Disproportionation reaction 32

33. Concentration of Solutions 33

34. Concentration How many moles of potassium dichromate are in 250 mL solution whose concentration is 2.16 M? Moles = Molarity X Volume Convert 250 mL to L: ◦ 0.250 L Plug numbers into the equation: ◦ Moles = (2.16 M) (0.250 L) ◦ Moles = 0.54 mol K 2 Cr 2 O 7 34

35. Concentration 35

36. Dilution Dilution: the procedure for preparing a less concentrated solution from a more concentrated one. M initial V initial = M final V final What volume of 8.61 M stock solution of H 2 SO 4 to prepare 5.00 X 10 2 mL of a 1.75 M H 2 SO 4 solution? (1.75 M)(5.00 X 10 2 mL) = (8.61M) V 2 V 2 = 102 mL 36

37. Gravimetric Analysis 37

38. Gravimetric Analysis 38

39. Acid-Base Titrations Titration: a solution of accurately known concentration. Standard solution: add gradually to another solution of unknown concentration until the chemical reaction between the two solutions is complete. Equivalence point: the point at which an acid has completely reacted with or been neutralized by a base. 39

40. Acid-Base Titrations 40

41. Redox Titrations 41