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1 Chemical Reactions Chapter 4, part 2: Reactions in Aqueous Solution.

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Presentation on theme: "1 Chemical Reactions Chapter 4, part 2: Reactions in Aqueous Solution."— Presentation transcript:

1 1 Chemical Reactions Chapter 4, part 2: Reactions in Aqueous Solution

2 2 The dissolving process When a solid is put into a liquid, solute-solute attractions compete with solute-solvent & solvent-solvent attractions

3 3 Solubility Lets assume our solvent is water... If solute-water attractions > solute-solute & water-water attractions, solute particles are pulled out one by one into the water: The solute is SOLUBLE in water

4 4 Solubility But if solute-water attractions < solute-solute & water-water attractions, solute particles remain together: The solute is INSOLUBLE in water

5 5 Solution conductivity Solution conductivity depends on type of solute particles

6 6 Solution conductivity Ionic solutes (salts) made of cations & anions Ions DISSOCIATE (separate) during dissolving Molecular solutes made of molecules each solute particle that moves into solution is identical

7 7 Solution conductivity Solutions of ionic solutes contain independent mobile ions Solution conducts electricity Solute is an ELECTROLYTE Solute also conducts when melted, but not when solid (ions cant move)

8 8 Solution conductivity Solutions of most molecular solutes contain independent neutral molecules Solution does not conduct electricity Solute is a NONELECTROLYTE Solute also does not conduct when melted or solid

9 9 Solution conductivity Some molecular compounds (acids & bases) react with water to produce ions, as if they dissociated A few acids & bases do this very well, producing lots of ions = strong electrolytes, strong acids & bases Most acids & bases do this weakly, producing a few ions = weak electrolytes, weak acids & bases

10 10 Strong acids & bases HClHNO 3 LiOHMg(OH) 2 HBrH 2 SO 4 NaOHCa(OH) 2 HIHClO 4 KOHSr(OH) 2 etc.Ba(OH) 2 STRONG ACIDS STRONG BASES All other acids & bases are WEAK

11 11

12 12 Predicting Electrolytes All soluble salts and strong acids & bases are strong electrolytes Weak acids & bases are weak electrolytes All other molecular compounds are nonelectrolytes

13 13 Precipitation reactions

14 14 Precipitation reactions spectator ions precipitate

15 15 Predicting Precipitation Reactions To predict whether a precipitate will form, you need to know which compounds are soluble (no ppt) and which are insoluble (ppt forms) Memorize the guidelines in Table 4.1 on page 154 in your text! Add this guideline: CrO 4 2– acts like SO 4 2–

16 16 Solubility guidelines: soluble Compounds of these ions are generally soluble and do NOT form precipitates: Alkali metals (group 1A, except Li 1+ ) Ammonium (NH 4 1+ ) Nitrates (NO 3 1– ) and acetates (C 2 H 3 O 2 1– ) Chlorides, bromides, iodides, except Pb 2+, Ag 1+, Hg 2 2+ Sulfates (SO 4 2– ) & chromates (CrO 4 2– ), except Sr 2+, Ba 2+, Pb 2+, and Hg 2 2+

17 17 Solubility guidelines: insoluble Compounds of these ions are generally insoluble and DO form precipitates: Hydroxides (OH 1– ) and Sufides (S 2– ) Alkali metals (group 1A) and ammonium (NH 4 1+ ) are soluble Sulfides of group 2A metals are generally soluble Hydroxides of Ca 2+, Sr 2+, and Ba 2+ are soluble Carbonates (CO 3 2– ) and phosphates (PO 4 3– )

18 18 Net ionic equations AgNO 3 (aq) + NaCl (aq) AgNO 3 (aq) + NaCl (aq) AgCl + NaNO 3 AgNO 3 (aq) + NaCl (aq) AgCl (s) + NaNO 3 (aq) The (aq) substances are dissociated: Ag 1+ (aq) + NO 3 1– (aq) + Na 1+ (aq) + Cl 1– (aq) AgCl (s) + Na 1+ (aq) + NO 3 1– (aq)

19 19 Net ionic equations Na 1+ and NO 3 1– are spectator ions Ag 1+ (aq) + NO 3 1– (aq) + Na 1+ (aq) + Cl 1– (aq) AgCl (s) + Na 1+ (aq) + NO 3 1– (aq) The net ionic equation omits spectators: Ag 1+ (aq) + Cl 1– (aq) AgCl (s)

20 20 Is that really a spectator? An ion is a spectator if and only if it is in exactly the same form in the products and reactants: Na 2 CO 3 (aq) + BaCl 2 (aq) BaCO 3 (s) + 2 NaCl (aq) CO 3 2– (aq) + Ba 2+ (aq) BaCO 3 (s) Na 2 CO 3 (s) + 2 HCl (aq) 2 NaCl (aq) + H 2 O + CO 2 (g) Na 2 CO 3 (s) + 2 H 1+ (aq) 2 Na 1+ (aq) + H 2 O + CO 2 (g) Only Cl 1– is a spectator Na 1+ is not a spectator because it was (s), then (aq)

21 21 Examples Indicate whether a ppt forms and if so, complete the reaction as a balanced net ionic equation: AlCl 3 (aq) + KOH (aq) K 2 SO 4 (aq) + FeBr 3 (aq) CaI 2 (aq) + Pb(NO 3 ) 2 (aq) Na 3 PO 4 (aq) + AlCl 3 (aq) Al 2 (SO 4 ) 3 (aq) + BaCl 2 (aq) (NH 4 ) 2 CO 3 (aq) + Pb(NO 3 ) 2 (aq)

22 22 Acids Acids produce H 3 O 1+ in aqueous solution: Strong acids are molecular compounds that react completely with water to produce H 3 O 1+ : HCl (g) + H 2 O H 3 O 1+ (aq) + Cl 1– (aq) For convenience, we often show acids as simply dissociating to produce H 1+ : HCl (g) H 1+ (aq) + Cl 1– (aq) There are only 6 strong acids (memorize them, pg 161)

23 23 Acids Acids produce H 3 O 1+ in aqueous solution: Weak acids are molecular compounds that react incompletely with water: HC 2 H 3 O 2 (aq) + H 2 O H 3 O 1+ (aq) + C 2 H 3 O 2 1– (aq) HC 2 H 3 O 2 (aq) H 1+ (aq) + C 2 H 3 O 2 1– (aq) All acids that are not strong are weak

24 24 Acids H 1+ is a proton The reaction with water is proton transfer: HCl (g) + H 2 O H 3 O 1+ (aq) + Cl 1– (aq) HC 2 H 3 O 2 (aq) + H 2 O H 3 O 1+ (aq) + C 2 H 3 O 2 1– (aq) Acids with one H 1+ to transfer are monoprotic acids HCl HNO 3 HC 2 H 3 O 2 Acids with more than one H 1+ to transfer are polyprotic acids H 2 SO 4 H 3 PO 4 H 2 C 3 H 2 O 4

25 25 Acids Polyprotic acids produce H 3 O 1+ in steps: H 2 SO 4 (aq) + H 2 O H 3 O 1+ (aq) + HSO 4 1– (aq) HSO 4 1– (aq) + H 2 O H 3 O 1+ (aq) + SO 4 2– (aq) For H 2 SO 4 the first step is strong and the second weak H 2 C 2 O 4 (aq) + H 2 O H 3 O 1+ (aq) + HC 2 O 4 1– (aq) HC 2 O 4 1– (aq) + H 2 O H 3 O 1+ (aq) + C 2 O 4 2– (aq) For all other polyprotic acids, the first step is weak and the second step is weaker

26 26 Bases Bases produce OH 1– in aqueous solution: Strong bases are ionic hydroxide compounds that are completely dissociated in water: NaOH (s) Na 1+ (aq) + OH 1– (aq) The strong bases are the hydroxides of group 1A and 2A metals (memorize them) Weak bases are molecular compounds that react incompletely with water to produce OH 1– : NH 3 (aq) + H 2 O (l) NH 4 1+ (aq) + OH 1– (aq) Amines (–NH 2 ) are weak bases

27 27 Neutralization Acids and bases neutralize each other The H 1+ from the acid is transferred to the base: HCl (aq) + NaOH (aq) HOH (l) + NaCl (aq) H 1+ (aq) + OH 1– (aq) HOH (l) The base is not always an OH 1– compound: HCl (aq) + NH 3 (aq) NH 4 1+ (aq) + Cl 1– (aq) H 1+ (aq) + NH 3 (aq) NH 4 1+ (aq)

28 28 Neutralization and net ionic equations It is important to recognize strong acids and bases when writing net ionic equations HCl (aq) + NaOH (aq) H 2 O (l) + NaCl (aq) H 1+ (aq) + OH 1– (aq) H 2 O (l) HC 2 H 3 O 2 (aq) + NaOH (aq) H 2 O (l) + NaC 2 H 3 O 2 (aq) HC 2 H 3 O 2 (aq) + OH 1– (aq) H 2 O (l) + C 2 H 3 O 2 1– (aq) The weak acid is not significantly ionized, so acetate is not a spectator (even if it is aq)

29 29 Examples Write the molecular and net ionic equations for HNO 3 + NaOH HF + KOH HC 2 H 3 O 2 + NH 3 H 2 SO 4 + Ba(OH) 2 H 2 C 2 O 4 + NaOH HCHO 2 + Ca(OH) 2

30 30 Gas forming reactions Some neutralizations produce a gas: CO 3 2– + 2 H 1+ H 2 CO 3 H 2 CO 3 is unstable and decomposes immediately H 2 CO 3 CO 2 (g) + H 2 O The overall reaction is CO 3 2– + 2 H 1+ CO 2 (g) + H 2 O

31 31 Gas forming reactions Memorize these gas-formers (pg 166): SO 3 2– + 2 H 1+ SO 2 (g) + H 2 O HSO 3 1– + H 1+ SO 2 (g) + H 2 O CO 3 2– + 2 H 1+ CO 2 (g) + H 2 O HCO 3 1– + H 1+ CO 2 (g) + H 2 O S 2– + 2 H 1+ H 2 S (g) NH OH 1– NH 3 (g) + H 2 O burning sulfur smell rotten egg smell ammonia

32 32 Redox Redox (oxidation-reduction) reactions are those in which electrons are transferred The loss of electrons is oxidation The gain of electrons is reduction LEO says GER

33 33 Oxidation states The oxidation state (O.S.) or oxidation number is a convenient but artificial way to describe the electron environment around an atom It is related to the number of electrons gained, lost, or apparently used in forming compounds Oxidation states are assigned using the rules on page 169 of your text (memorize these in order)

34 34 Assigning oxidation states 1. The O.S. of each atom in an element is zero 2. The O.S. of a monoatomic ion is equal to its charge 3. The total of the O.S. of all atoms in any species (formula unit, molecule or ion) equals the charge on that species 4. In compounds, metals always have a positive O.S. Group 1A metals are always O.S. +1 and Group 2A metals are always O.S For nonmetals in compounds, the O.S. of fluorine is –1. the O.S. of hydrogen is +1. the O.S. of oxygen is –2. 6. In binary compounds, the O.S. of a Group 7A element is –1, Group 6A element –2, and Group 5A element –3.

35 35 Examples What is the oxidation state of each element in S 8 Cr 2 O 7 2– Cl 2 O KO 2 What is the oxidation state of each element in S 2 O 3 2– Hg 2 Cl 2 KMnO 4 H 2 CO

36 36 Redox In a redox reaction, atoms change (O.S.) The element oxidized loses electrons and its O.S. becomes more positive The element reduced gains electrons and its O.S. becomes more negative

37 37 Examples Which of these reactions is a redox reaction? Identify the species oxidized and reduced. HCl (aq) + NaOH (aq) H 2 O + NaCl (aq) 2 Pb(NO 3 ) 2 (s) 2 PbO (s) + 4 NO 2 (g) + O 2 (g) NH 4 Cl (s) + NaOH (aq) NH 3 (g) + H 2 O + NaCl (aq) æ Identify the species oxidized and reduced: 5 VO 2+ (aq) + MnO 4 1– (aq) + H 2 O 5 VO 2 1+ (aq) + Mn 2+ (aq) + 2 H 1+ (aq)

38 38 Agents of Oxidation and Reduction An agent makes something happen An oxidizing agent makes oxidation happen by being reduced A reducing agent makes reduction happen by being oxidized The agent is the entire species in which the oxidized or reduced atom appears

39 39 Examples Identify the elements oxidized & reduced and the oxidizing & reducing agents in 2 NO 2 (g) + 7 H 2 (g) 2 NH 3 (g) + 4 H 2 O (g) 5 H 2 O 2 (aq) + 2 MnO 4 1– (aq) + 6 H 1+ (aq) 8 H 2 O + 2 Mn 2+ (aq) + 5 O 2 (g) S 2 O 3 2– (aq) + 4 Cl 2 (aq) + 5 H 2 O 2 HSO 4 1– (aq) + 8 H 1+ (aq) + 8 Cl 1– (aq) 6 Fe 2+ (aq) + 14 H 1+ (aq) + Cr 2 O 7 2– (aq) 6 Fe 3+ (aq) + 2 Cr 3+ (aq) + 7 H 2 O S 2 O 3 2– (aq) + 2 H 1+ (aq) S (s) + SO 2 (g) + H 2 O

40 40 Reactions You now know how to write & balance 4 types of reactions combustion precipitation acid-base neutralization gas-forming and how to recognize redox reactions

41 41 Reaction quizzes Reaction quizzes (RQ) will replace NQ I give you the names of the reactants You write the formulas of the reactant and the formulas of the products Cross out spectator ions Write the balanced net ionic equation You need not include state symbols such as (aq) or (g)

42 42 Example Solutions of silver nitrate and potassium chloride are mixed AgNO 3 (aq) + KCl (aq) AgNO 3 (aq) + KCl (aq) AgCl (s) + KNO 3 (aq) cross out K 1+ and NO 3 1– as spectators Ag 1+ + Cl 1– AgCl check: it is already balanced

43 43 Example Hydrochloric acid solution is added to solid sodium hydrogen carbonate HCl (aq) + NaHCO 3 (s) HCl (aq) + NaHCO 3 (s) H 2 O + CO 2 (g) + NaCl (aq) cross out Cl 1– as a spectator do NOT cross out Na 1+ because NaHCO 3 is solid H 1+ + NaHCO 3 H 2 O + CO 2 + Na 1+ check: it is already balanced

44 44 Example Ethoxyethane burns in air C 2 H 5 OC 2 H 5 + O 2 C 2 H 5 OC 2 H 5 + O 2 CO 2 + H 2 O there are no spectators in combustion balance C 2 H 5 OC 2 H O 2 4 CO H 2 O

45 45 Example Solutions of nitrous acid and sodium hydroxide are mixed HNO 2 (aq) + NaOH (aq) HNO 2 (aq) + NaOH (aq) H 2 O + NaNO 2 (aq) cross out Na 1+ as a spectator do NOT cross out NO 2 1– because HNO 2 is a weak acid and is not significantly dissociated in solution! HNO 2 + OH 1– H 2 O + NO 2 1– check: it is balanced

46 46 Titrations The text includes examples of acid-base titration and redox titration These are just stoichiometry problems, based on a particular type of reaction Do examples 5-9 and 5-10 to practice these They will be on the test, and you may have to write and/or balance your own equations

47 47 Exercise 26 Both NH 3 (aq) and HC 2 H 3 O 2 (aq) conduct weakly, but when these solutions are mixed, the resulting solution conducts electricity very well. Explain.

48 48 Exercise 32 Assuming volumes are additive, what is [NO 3 1– ] in a solution obtained by mixing 275 mL M KNO 3, 328 mL M Mg(NO 3 ) 2, and 784 mL H 2 O?

49 49 Exercise 34 Predict whether a reaction occurs, and if so, write a net ionic equation: a. AgNO 3 (aq) + CuCl 2 (aq) b. Na 2 S (aq) + FeCl 2 (aq) c. Na 2 CO 3 (aq) + AgNO 3 (aq)

50 50 Exercise 39 Write net ionic equations to represent these neutralizations of stomach acid (HCl): a. sodium bicarbonate b. calcium carbonate c. magnesium hydroxide

51 51 Exercise 42 Give net ionic equations for the preparation of each salt by acid-base neutralization: (NH 4 ) 2 HPO 4, NH 4 NO 3, (NH 4 ) 2 SO 4

52 52 Exercise 47 Balance these redox reactions in acid solution: a. MnO 4 1– + I 1– Mn 2+ + I 2 (s) b. BrO 3 1– + N 2 H 4 Br 1– + N 2 c. VO 4 3– + Fe 2+ U UO NO (g)

53 53 Exercise 49 Balance these redox reactions in basic solution: a. CN 1– + MnO 4 1– MnO 2 (s) + CNO 1– b. [Fe(CN) 6 ] 3– + N 2 H 4 [Fe(CN) 6 ] 4– + N 2 c. Fe(OH) 2 (s) + O 2 (g) (OH) 3 (s) C 2 H 5 OH + MnO 4 1– C 2 H 3 O 2 1– + MnO 2 (s)

54 54 Exercise 52 Write a balanced equation for each redox reaction: a. Oxidation of nitrite ion to nitrate ion by permanganate ion in acid solution, producing Mn 2+ b. Reduction of H 2 S with SO 2. Both are converted to elemental sulfur and water. c. Reaction of sodium metal with hydroiodic acid. d. Reduction of VO 2+ to V 3+ by zinc metal, in acidic solution.

55 55 Exercise 53 Identify the oxidizing and reducing agents in these reactions: a. 5 SO 3 2– + 2 MnO 4 1– + 6 H + 5 SO 4 2– + 2 Mn H 2 O b. 2 NO H 2 2 NH H 2 O c. 2 [Fe(CN) 6 ] 4– + H 2 O H + 2 [Fe(CN) 6 ] 3– + 2 H 2 O

56 56 Exercise 60 a. How many mL of concentrated HCl (d = 1.19 g/mL, 38% HCl by mass) must be diluted to 20.0 L to prepare 0.10 M HCl? b mL of the 0.1 M HCl requires mL of M NaOH for titration. What is the molarity of the HCl? c. Why could you not prepare M HCl simply by diluting concentrated HCl?

57 57 Exercise mL battery acid (H 2 SO 4 ) requires Ml M NaOH for complete neutralization. Is the concentration of the battery acid within the desired range (4.8 – 5.3 M)?

58 58 Exercise 67 Iron ore weighing g is dissolved in HCl to produce Fe 2+, which is then titrated with mL M K 2 Cr 2 O 7. What is the % Fe by mass in the ore? 6 Fe H + + Cr 2 O 7 2– 6 Fe Cr H 2 O

59 59 Half-reactions A half reaction describes either the oxidation or the reduction part of a redox reaction It includes the number of electrons lost or gained The overall reaction is the sum of two half-reactions, one oxidation and one reduction

60 60 Examples 5-5A & 5-5B Represent reaction 4.2 with half-reactions and the overall net reaction: 2 Al (s) + 6 HCl (aq) 2 AlCl 3 (aq) + 3 H 2 (g) Represent the reaction of chlorine gas with aqueous sodium bromide to produce liquid bromine and aqueous sodium chloride with half- reactions and a net equation

61 61 Balancing Redox Equations Redox equations require electrical balance as well as material (atom) balance Write balanced half reactions for oxidation & reduction First balance atoms, adding H 1+ and H 2 O as needed When atoms are balanced, add e – to balance charge. Multiply each half-reaction by a factor such that the number of electrons lost and gained are the same Combine the half reactions and cancel species that are identical on both sides

62 62 Examples 5-6A & 5-6B Balance these equations in acid solution Fe 2+ (aq) + MnO 4 1– (aq) Fe 3+ (aq) + Mn 2+ (aq) UO 2+ (aq) + Cr 2 O 7 2– (aq) UO 2 2+ (aq) + Cr 3+ (aq)

63 63 Redox in basic solution To balance a redox equation in basic solution, Balance it as if it is in acid solution After combining the half-reactions, add equal moles of OH 1– to both sides such that all H 1+ is neutralized to form H 2 O Simplify

64 64 Examples 5-7A & 5-7B Balance these equations in basic solution S (s) + OCl 1– (aq) SO 3 2– (aq) + Cl 1– (aq) MnO 4 1– (aq) + SO 3 2– (aq) MnO 2 (s) + SO 4 2– (aq)

65 65 Disproportionation In disproportionation reactions, the same element is both oxidized and reduced 2 H 2 O 2 (aq) 2 H 2 O + O 2 (g) S 2 O 3 2– (aq) S (s) + SO 2 (g)

66 66 Common Redox agents Common oxidizing agents include MnO 4 1– (permanganate) in acid solution Cr 2 O 7 2– (dichromate) in acid solution O 2 or O 3 Common reducing agents include S 2 O 3 2– (thiosulfate) H 2


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