Section 4: Aqueous Reactions

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1 Section 4: Aqueous Reactions 1. Solution composition 2. Electrolytes and nonelectrolytes 3. Acids, bases, and salts 4. Neutralization ti reactions 5. Precipitation reactions 6. Oxidation/reduction reactions 7. Molarity 1

2 Water has remarkable chemical and physical properties High melting and boiling points Expands when freezes (ice floats) Dissolves a wide variety of substances (the universal solvent) Provides an ideal environment in which numerous chemical reactions can occur (aqueous reactions) The chemical reactions of life require and The chemical reactions of life require, and occur in, water 2

3 Solutions, solvents, and solutes Solution: a homogeneous mixture of 2 or more substances Solvent: the substance in the solution present in greatest quantity Solutes: other substances that are dissolved in the solvent Solution concentration: the amount of solute dissolved in a given quantity of solvent or solution An aqueous solution is a solution where water is the solvent 3

4 Some common aqueous solutions Solutions of sugars (glucose, fructose, sucrose, lactose, etc) Solutions of alcohols (methanol, ethanol, etc) Solutions of salts (sodium chloride, potassium chloride, ammonium sulfate, etc) Solutions of acids Solutions of bases 4

5 The solution process 5

6 The solution process 1. Break intermolecular attractions in pure solute to free up individual id molecules l (solute) n n solute 2. Disrupt water-water intermolecular attractions to provide room for solute to dissolve (water) m m water 3. Form new solute-water intermolecular forces n solute + m H 2 O (solute) n (water) m 6

7 Electrolytes and nonelectrolytes Electrolytes: compounds that dissolve in water to produce ions (soluble ionic compounds, acids and bases in water) dissociate or ionize in solution Nonelectrolytes: compounds that dissolve in water, and yet do not form ions in solution (soluble molecular compounds in water) dissolve but don t ionize 7

8 Sodium chloride dissolves in water to form hydrated sodium and choride ions + - 8

9 Electrolytes: acids, bases, and salts Acids: Hydrogen-containing compounds that dissociate in water to produce hydrogen ion and a partner anion Compounds that t increase [H + ] in aq. solutions Bases: compounds that react with (accept) hydrogen ions, producing water Compounds that increase [OH-] in aq. solutions Salts: Ionic compounds that are neither acids or bases 9

10 Acids: Hydrogen-containing compounds that dissociate in water to produce hydrogen ion (H + ) and a partner anion Strong acids: dissociate (ionize) completely in solution Weak acids: dissolve completely, but only partially ionize in solution Polyprotic acids: have more than one ionizable proton 10

11 Bases: (1) compounds that react with (accept) hydrogen ions, producing water. (2) Compounds that increase [OH - ] in aq. solutions. Metal hydroxides Ammonia Strong bases: dissociate (ionize) completely in solution Weak bases: Ionize only partially in solution, or dissolve incompletely in solution 11

12 Strong acids and bases Common strong acids HCl, HBr, HI (but not HF) HClO 4 HNO 3 Common strong bases Group 1A metal hydroxides (LiOH, NaOH, KOH, RbOH, CsOH). These dissociate completely at 1 M and higher concentrations H 2 SO 4 (only the first H + ionizes completely) The heavier Group 2A metal hydroxides dissociate completely but only at low concentrations 12

13 Salts: Compounds that dissociate (ionize) in water but are not acids or bases 13

14 Electrolytes Strong electrolytes: salts, acids, and bases that dissolve completely or nearly completely in water to form ions Weak electrolytes: l t Compounds that t dissolve in water, but produce only a small concentration of ions when dissolved (partial dissociation; partial ionization) 14

15 Electrolytes and nonelectrolytes distilled water: a nonelectrolyte tap water: a weak electrolyte 15

16 Electrochemistry is a fascinating area The chemistry going on in this simple demonstration could be the solution to much of of our future energy needs 16

17 Solubility: a term that refers to how much of a solute can dissolve in a given amount of solvent (grams solute dissolved/100 ml solution) Ionic compounds Molecular compounds Acids and bases 17

18 Degrees of solubility soluble in water: glucose, NaCl, NaOH, acetic acid, MgSO 4 ( soluble means a significant amount can dissolve) Insoluble in water: CaCO 3, AgCl, C 6 H 14, I 2 ( insoluble means a very little amount dissolves) miscible with water: ethanol, glycerol, formaldehyde ( miscible means soluble at all proportions) 18

19 Solubility of salts Most salts are classified as soluble Some salts are insoluble (poorly soluble) The distinction between soluble and insoluble is based on whether a significant amount dissolves in water (see next slide) Don t memorize which salts are soluble and insoluble at this point, but some being familiar with some general guidelines will be useful for you 19

20 Salt Solubility (g/100 g water at 25 C Classification NaCl 35.9 Largely soluble NaNO Largely soluble Ca(CH 3 COO) 2 ~33 Largely soluble MgSO 4 ~25 Largely soluble CaSO 4 ~0.2 Largely insoluble ACl AgCl Largely insoluble CaF 2 ~0.002 Largely insoluble NH 4 F 45 Largely soluble PbCl 2 ~4.5 Slightly soluble Pb(NO 3 ) 2 ~60 Largely soluble CaCO Largely insoluble Na 2 CO 3 ~230 Largely soluble PbS 2 ~2.6 x 10 9 Largely insoluble Pb(CH 3 COO) 2 ~50 Largely soluble 20

21 Solubility of some commonly encountered salts in water All ionic compounds with nitrate (NO - 3 ), group 1A metal ions, and ammonium ion (NH 4+ ) are soluble All binary salts of Cl -, Br -, I - are soluble, except with Ag +, Hg +, Pb 2+ Nearly all acetate (CH 3 COO - ) salts are soluble Most sulfate (SO 2-4 ) salts are soluble (exceptions: BaSO 4 4, CaSO 4, PbSO 4, SrSO 4 Most hydroxide (OH - ) salts are insoluble (ex. w/group 1A metal ions or NH + 4, heavier group 2A hydroxides are slightly soluble) Most oxide (O 2- ) and sulfide (S 2- )salts are insoluble (ex. w/group 1A metal ions or NH 4+ ) Most phosphate and carbonate salts are insoluble (ex. w/group 1A metal ions or NH 4+ ) Many F - salts are insoluble (ex. w/group 1A metal ions or NH 4+ and some others) Blah, blah, blah.. That s enough for now 21

22 Summary. Many salts are soluble and hence strong electrolytes Most acids are weak acids and hence weak electrolytes Group 1A and the heavier group 2A metal hydroxides are strong bases and hence strong electrolytes Ammonia is a weak base and hence a weak electrolyte Soluble molecular compounds that don t ionize in solution are nonelectrolytes 22

23 Common reactions in aqueous Precipitation Neutralization solutions Gas generation Metals: oxidation/reduction 23

24 Precipitation Reactions: Formation of an insoluble salt from the mixing of aqueous solutions of soluble salts Example: Silver nitrate (AgNO 3 ) is soluble in water. Sodium chloride (NaCl) is soluble in water. Silver chloride (AgCl) is insoluble in water. What will happen if solutions of AgNO 3 and NaCl are mixed together? Write a balanced chemical equation for the reaction. 24

25 Precipitation Reactions: Formation of an insoluble salt from the mixing of aqueous solutions of soluble salts Example: Silver nitrate (AgNO 3 ) is soluble in water. Sodium chloride (NaCl) is soluble in water. Silver chloride (AgCl) is insoluble in water. What will happen if solutions of AgNO 3 and NaCl are mixed together? Write a balanced chemical equation for the reaction. 25

26 More solution chemistry terms Molecular equation: shows the complete chemical formulas of reactants and products, not necessarily as they actually exist in solution Complete ionic equation: all soluble salts are shown as ions- i.e. in the form they actually exist in solutions Net ionic equation: includes only ions and molecules that are directly involved in a chemical reaction occurring in solution Spectator ions: ions in solution that are not involved in a chemical reaction, and appear in identical forms on both sides of the chemical equation 26

Write molecular, complete ionic, and net

iodide precipitates when potassium iodide

27 Write molecular, complete ionic, and net ionic equations for the formation of a precipitate of PbI 2 after mixing solutions of Pb(NO 3 ) 2 and KI Lead (II) iodide precipitates when potassium iodide is mixed with lead (II) nitrate. Author PRHaney licensed under the Creative Commons Attribution-Share Alike 3.0 Unported license. 27

28 Neutralization reactions of acids and bases A nonelectrolyte (water) and a salt or weak electrolyte is formed A reactant may be soluble or insoluble to begin with Sometimes a product is a gas, or reacts further to produce a gas Reactions can be represented by molecular, complete ionic, or net ionic equations 28

29 Neutralization reactions of acids and bases A nonelectrolyte (water) and a salt or weak electrolyte is formed Reactions can be represented by molecular, complete ionic, or net ionic equations Hydrochloric acid + sodium hydroxide Sulfuric acid + potassium hydroxide 29

30 Neutralization reactions of acids and bases A nonelectrolyte (water) and a salt or weak electrolyte is formed A reactant may be insoluble to begin with Insoluble hydroxide plus strong acid: Insoluble metal oxide + strong acid: 30

31 Neutralization reactions of acids and bases A nonelectrolyte (water) and a salt or weak electrolyte is formed Sometimes a product is a gas, or the product reacts further to produce a gas Strong acid + sodium sulfide or sodium cyanide Strong acid + carbonate or sulfite salt 31

32 Neutralization reactions of acids and bases A nonelectrolyte (water) and a salt or weak electrolyte is formed Sometimes a product is a gas, or the product reacts further to produce a gas Alka seltzer Baking powder 32

33 Common reactions in aqueous Precipitation Neutralization Gas generation Oxidation/reduction solutions 33

34 Oxidation/reduction reactions: chemical reactions where electrons are transferred from one reactant to another 34

35 Oxidation/reduction reactions: chemical reactions where electrons are transferred from one reactant to another Na: 1s 2 2s 2 2p 6 3s 1 Cl: 1s 2 2s 2 2p 6 3s 2 3p 5 n = 3 n = 3 n = 1 n = 2 11p+ 17 p+ n = 1 n = 2 35

36 Oxidation/reduction reactions: chemical reactions where electrons are transferred from one reactant to another Na + : 1s 2 2s 2 2p 6 Cl - : 1s 2 2s 2 2p 6 3s 2 3p 6 n = n = 2 11p+ 17 p+ n = 1 n = 2 n = 3 36

37 Oxidation/reduction reactions Oxidation: lose electrons Metals lose electrons to obtain a Noble gas e- configuration (Li Li + + e - ) Reduction: gain electrons nonmetals often gain electrons to obtain a Noble gas electron configuration (Cl 2 + 2e - 2Cl - ) Some cations will gain electrons to form a stable reduced product (2H + + 2e - H 2 ) An oxidation reaction must be linked to a reduction reaction Oxidation numbers assigned to atoms and ions in compounds are used to keep track of electrons gained or lost during oxidation/reduction reactions 37

38 Oxidation numbers An atom in its elemental state has an oxidation number of zero, meaning it has neither gained nor lost electrons Oxidation numbers of H, N, O, C, and S in H 2, N 2, O 2, C (graphite), and S 8 are all zero For monatomic ions, the oxidation number is the same as the ion charge Monatomic metal cation Oxidation number Monatomic nonmetal anion Oxidation number Li + Na +, K +, Cu + plus one (+1) F - Cl -, Br -, I - minus one (-1) Mg 2+, Ca 2+, Sr 2+, Fe 2+ plus two (+2) O 2-, S 2- Se 2-, Te 2- minus two (-2) Al 3+, Fe 3+, Cr 3+, Ni 3+ plus three (+3) N 3-, P 3- minus three (-3) Note H, a nonmetal, can have an oxidation number of +1 in the hydrogen ion (H + ) or -1 in the hydride ion (H - ) The oxidation numbers of atoms in molecular compounds and polyatomic ions are assigned by rules we will discuss later 38

39 Metals lose electrons to obtain a Noble gas e- configuration (Li Li + + e - ) Nonmetals gain electrons to obtain a Noble gas electron configuration (Cl 2 + 2e - 2Cl - ) Some cations will gain electrons to form a stable reduced product (2H + +2e - H 2 ) Alkali metal plus halogen gas: Alkali metal plus water: Metal plus strong acid: 39

40 Solution concentration The amount of solute dissolved in a given quantity of solvent or solution 100L particles/liter 25 particles/liter 50 particles/liter 40

41 Molarity Moles solute Molarity = M = Volume of solution (liters) mols L # of particles Defined volume [solute] 41

42 (1) Add proper mass of solute to prepare defined volume of solution, add some water (2) Mix to (3) Add more (4) Solution of dissolve solute water to reach defined molarity final desired volume 42

43 Preparing a 1 M solution of sodium chloride Na Cl FW NaCl = amu = g/mol 43

44 Preparing a 1 M solution of sodium chloride FW NaCl = g/mol NaCl g NaCl 1.00 L 1 mol/l NaCl 44

45 Interconversion of mass, mols, and number of particles Molar Avogadro s mass mass moles number molecules Interconversion of mass, mols, and concentration of particles Molar Volume of mass mass moles solution molarity mass solution mols L 45

46 Potassium permanganate (KMnO 4 ), a common lab chemical, is a purple-black solid that dissolves in water to give a deep purple solution. If you dissolve g of KMnO 4 in enough water to give 250 ml of solution, what is [KMnO 4 ] in mols/l? FW KMnO4 = g/mol 46

47 If you dissolve g of KMnO 4 in enough water to give 250 ml of solution, what is [KMnO 4 4] in mols/l? Molar Volume of mass mass moles solution molarity FW KMnO4 = g/mol V = 0.25 l 47

48 Molarity as a conversion factor Molar Volume of mass mass moles solution molarity Molar Molarity of mass mass moles solution volume g mol mol l 48

49 How many mols of HCl are present in 250 ml of 12 M HCl? Molar Molarity of mass mass moles solution volume g mol mol l 49

50 You need 5.0 mols of HCl. What volume of 12 M HCl would provide these 5.0 mols? Molar Molarity of mass mass moles solution volume g mol mol l 50

51 For a certain experiment, you need 250 ml of 1.50 M sodium carbonate (Na 2 CO 3 ). How do you prepare p this solution? Molar Molarity of mass mass moles solution volume Molar Volume of mass mass moles solution molarity 51

52 For a certain experiment, you need 250 ml of 1.50 M sodium carbonate (Na 2 CO 3 ). How do you prepare this solution? Molar Molarity of mass mass solution moles volume 52

53 Changing solution concentration by dilution 20 particles/100 ml 20 particles/200 ml 53

54 Dilution What is the concentration of an NaOH solution prepared by diluting 0.25 L of 6.0 M NaOH to a final volume of 1.25 L by addition of water? 6.0 mol NaOH L L 1.25 L V 1 V M 1 moles 2 M 2 54

55 Dilution What volumes of f60mn 6.0 NaOH and water are required to prepare 0.50 liters of 1 M NaOH? 0.50 L 1.0 mol NaOH 1.0 L 1.0 L 6.0 mol NaOH M 1 M V 1 moles 2 V 2 55

56 M 1 xv 1 = moles M 2 x V 2 = moles 2 2 moles = moles So: Dilution VV 1 VV M1 moles 2 M V 2 1 V 2 M 1 MM l 2 1 V 1 moles V 2 M 1 M 2 M 1 x V 1 = M 2 x V 2 56

57 Review some solution terms Molecular l equation: shows the complete chemical lformulas of reactants and products, not necessarily as they actually exist in solution Complete ionic equation: all soluble salts are shown as ions- i.e. in the form they actually exist in solutions Net ionic equation: includes only ions and molecules that are directly involved in a chemical reaction occurring in solution Spectator ions: ions in solution that are not involved in a chemical reaction, and appear in identical forms on both sides of the chemical equation 57

58 What concentrations of species are present in solutions of strong electrolytes? Molecular equation: shows the complete chemical formulas of reactants and products, not M NaCl M Na necessarily as they actually 2 SO 4 exist in solution Complete ionic equation: all soluble salts are shown as ions- i.e. in the form they actually exist in solutions Net ionic equation: includes only ions and molecules that are directly involved in a chemical reaction occurring in solution Spectator ions: ions in solution that are not involved in a chemical reaction, and appear in identical forms on both sides of the chemical equation 58

59 Useful problem solving strategies for chemical reactions occurring in aqueous solutions Interconverting moles, molarity, and volume: M V V moles M M Dilution problems: M1 V 1 V moles 2 V 1 V 2 M 2 1 V M 1 M V 1 moles 2 V 2 M 1 M 2 Interconverting mass, moles, and volume in aqueous chemical reactions: Molar mass M a a or V a mass a moles a M a or V a Stoichiometric relationship in a reaction Molar mass mass b M b moles b or V b b M b or V b 59

60 Problem solving strategy for chemical reactions occurring in aqueous solutions Molar mass M or V a a a mass a moles a M a or V a Molar mass b Stoichiometric relationship in a reaction M b or V b massbb moles b M b or V b 60

61 Sodium bicarbonate (NaHCO 3 ) is often used to neutralize spills of acids such as sulfuric acid (H 2 SO 4 ). 2NaHCO 3(s) + H 2 SO 4(aq) Na 2 SO 4(aq) + 2H 2 O (l) + 2CO 2(g) What mass of NaHCO 3 would be required to neutralize 25 ml of 6.0 M H 2 SO 4? Molar mass a M a or V a mass a moles a M a or V a Stoichiometric relationship in a reaction Start here Molar mass b M b or V b b b b massb moles b M b or V b End here 61

62 2NaHCO 3(s) + H 2 SO 4(aq) Na 2 SO 4(aq) + 2H 2 O (l) + 2CO 2(g) What mass of NaHCO 3 would be required to neutralize 25 ml of 6.0 M H 2 SO 4? Molar mass mass a M a or V a a moles a M a or V a Stoichiometric relationship in a reaction Start here mass b Molar mass b moles b M b or V b M b or V b End here 62

63 What volume of M HClO 4 would be required to neutralize 50.0 ml of M NaOH? Molar mass a M a or V a mass a moles a M a or V a Stoichiometric relationship in a reaction Start here Molar mass mass b M b or V b b moles b M b or V b End here 63

64 What volume of M HClO 4 would be required to neutralize 50.0 ml of M NaOH? 64

65 What volume of M HCl would be required to neutralize 2.87 g of Mg(OH) 2? 65

66 What is the total concentration of all solute species present in a solution prepared by mixing 50.0 ml of 0.20 M KClO 3 with 25.0 ml of 0.20 M Na 2 SO 4?? 66

Solubility Rules See also Table 4.1 in text and Appendix G in Lab Manual

Solubility Rules See also Table 4.1 in text and Appendix G in Lab Manual Ch 4 Chemical Reactions Ionic Theory of Solutions - Ionic substances produce freely moving ions when dissolved in water, and the ions carry electric current. (S. Arrhenius, 1884) - An electrolyte is a