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Chemistry II Aqueous Reactions and Solution Chemistry Chapter 4.

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1 Chemistry II Aqueous Reactions and Solution Chemistry Chapter 4

2 In this chapter we will consider chemical processes that occur in aqueous solutions: precipitation reactions, acid base reactions and oxidation – reduction reactions.

3 We will then consider concentrations and how the concepts of stoichiometry can be applied to concentrations.

4 Water has many properties that allow it to help support life. One of these properties is that it can dissolve a wide variety of materials. For this reason water is often referred to as what?

5 The universal solvent

6 What are aqueous solutions?

7 Solutions in which water is the dissolving medium.

8 Section 1: General Properties of Aqueous Solutions

9 Define solution:

10 A homogeneous mixture of two or more substances.

11 What is the difference between solvent and solute?

12 The substance that is present in greater quantity is usually called the solvent. The other substances in the solution are solutes. Solutes are dissolved in the solvent.

13 Electrolytic Properties Pure water is not a good conductor of electricity. The presence of ions causes aqueous solutions to become good conductors. Ions carry the charge from one electrode to the next.

14

15 What is an electrolyte?

16 A substance whose aqueous solutions contains ions.

17 What is a non-electrolyte?

18 A substance that does not form ions in solution.

19 Ionic Compounds in water

20 What does it mean when we say an ionic solid dissociates into its component ions as it dissolves?

21 Each ion separates from the solid structure and disperses throughout the solution.

22

23 What is a polar molecule. Explain the significance of this fact with respect to the dissociation of ionic solids.

24

25 A polar molecule has one end that has a partial positive charge and a partial negative charge.

26 Although water is an electrically neutral molecule, one end of the molecule is electron rich and carries a partial negative charge. The hydrogen side of the molecule has a partial positive charge.

27 When ionic compounds dissolve the anions are surrounded by the water molecules so that the hydrogen side of the molecule surrounds the anion. The cations are surrounded by the oxygen side of the water molecule. This configuration stabilizes the ions in solution.

28 How can we predict the charges of the ions present in solution?

29 Remember the formulas and charges of the common ions. i.e., Na 2 SO 4 will separate into two Na + ions and one SO 4 2- ions. In solution for every one sodium sulfate three ions are formed.

30 Molecular Compounds in Water

31 When a molecular compound dissolves in water, the solution usually consists of intact molecules dispersed throughout the solution. Usually molecular compounds are non- electrolytes. An exception to this rule is acids

32

33 In what way do acids appear to not follow the rule?

34 Acids are molecular compounds that will disassociate or ionize into ions in aqueous solutions. HCl H + + Cl -

35 The two categories of electrolytes are strong and weak electrolytes

36 What are strong electrolytes?

37 Those solutes that exist in solution completely or nearly completely as ions. Most ionic compounds and some acids and bases are strong electrolytes.

38 What are weak electrolytes? Those solutes that exist in solution in the form of molecules but only partially disassociated into ions.

39 Can you determine if a solute is a strong or weak electrolyte by how well it dissolves?

40 No, for example acetic acid (vinegar) is very soluble in water, but only partially dissociates into ions.

41 How can we indicate that an electrolyte is a weak electrolyte?

42 We can use double arrows to show that the reaction is significant in both directions. HC 2 H 3 O 2(aq) H + (aq) + C 2 H 3 O 2 - (aq) * The state of equilibrium between molecules and ions varies from one weak electrolyte to another.

43 How do chemists indicate the ionization of strong electrolytes?

44 With the use of a single arrow. HCl (aq) H + (aq) + Cl - (aq) The single arrow indicates that the ions have no tendency to recombine to molecules.

45 In a the next few sections we will learn how to predict if a compound is a strong electrolyte, weak electrolyte or non-electrolyte. For now, in general, soluble ionic compounds are always strong electrolytes.

46 How can we identify compounds are being ionic?

47 Ionic compounds are composed of metals and nonmetals NaCl FeSO 4 Al(NO 3 ) 3 NH 4 Br

48 The diagram on the right represents an aqueous solution of one of the following compounds: MgCl 2, KCl, or K 2 SO 4. Which solution does the drawing best represent?

49 If you were to draw diagrams (such as that shown on exercise 4.1) representing aqueous solutions of each of the following ionic compounds, how many anions would you show if the diagram contained six cations? (a) NiSO 4, (b) Ca(NO 3 ) 2, (c) Na 3 PO 4, (d) Al 2 (SO 4 ) 3

50 Homework Page

51 Section 2_ Precipitation Reactions

52 What are precipitation reactions?

53 Reactions that result in the formation of an insoluble product.

54 What is a precipitate?

55 An insoluble solid formed by a reaction in solution.

56 When do precipitation reaction occur? Precipitation reactions occur when certain pairs of oppositely charged ions attract each other so strongly that they form an insoluble ionic solid.

57 For example Pb(NO 3 ) 2(aq) + 2KI (aq) PbI 2(s) + 2KNO 3(aq)

58 Solubility Guidelines for Ionic compounds

59 What is solubility ?

60 The amount of substance that can be dissolved in a given quantity of solvent. Any substance with a solubility less than mol/L is referred to as insoluble.

61 Experimental observations have led to guidelines for predicting solubility (page 118 Table 4.1)

62 Sample exercise 4.2 Classify the following ionic compounds as soluble or insoluble in water: (a) sodium carbonate (Na 2 CO 3 ), (b) lead sulfate (PbSO 4 ).

63 Classify the following compounds as soluble or insoluble in water: (a) cobalt(II) hydroxide, (b) barium nitrate, (c) ammonium phosphate.

64 How can we predict whether a precipitate forms when we mix aqueous solutions of two strong electrolytes? 1.) note the ions present in the reactants.

65 How can we predict whether a precipitate forms when we mix aqueous solutions of two strong electrolytes? 1.) note the ions present in the reactants. 2.) Consider possible combinations of cations and anions.

66 How can we predict whether a precipitate forms when we mix aqueous solutions of two strong electrolytes? 3.) Use solubility guidelines to determine if any combinations are insoluble

67 How can we predict whether a precipitate forms when we mix aqueous solutions of two strong electrolytes? For example, will a precipitate form when Mg(NO 3 ) 2 and NaOH are mixed?

68 Precipitation reactions are a type of double replacement reactions. They are also known as exchange or metathesis reactions.

69 AX + BY AY + BX AgNO 3 + KCl AgCl + KNO 3

70 Exchange (Metathesis) Reactions In exchange reactions the chemical formulas of the products are based on the charges of the ions.

71 Sample exercise 4.3 (a) Predict the identity of the precipitate that forms when solutions of BaCl 2 and K 2 SO 4 are mixed. (b) Write the balanced chemical equation for the reaction.

72 (a) What compound precipitates when solutions of Fe 2 (SO 4 ) 3 and LiOH are mixed? (b) Write a balanced equation for the reaction.

73 Will a precipitate form when solutions of Ba(NO 3 ) 2 and KOH are mixed?

74 Ionic Equations

75 In writing chemical equations for reactions in aqueous solutions, it is often helpful to know if the dissolved substances are present mainly as molecules or as ions. For example: Molecular Equation- Pb(NO 3 ) 2(aq) + 2KI (aq) PbI 2(s) + 2KNO 3(aq)

76 Ionic Equation Pb +2 (aq) +2NO 3(aq) - +2K + (aq) + 2I - (aq) PbI 2(s) + 2K + (aq) + 2NO 3 - (aq)

77 What is a complete ionic equation? An equation such as the one above written with all strong electrolytes written as ions.

78 What are spectator ions? Ions that appear in identical forms on both the reactant and product side of the equation.

79 What are spectator ions? Pb +2 (aq) +2NO 3(aq) - +2K + (aq) + 2I - (aq) PbI 2(s) + 2K + (aq) + 2NO 3 - (aq)

80 What is a net ionic equation?

81 When the spectator ions are cancelled out we are left with the net ionic equation.

82 Pb +2 (aq) + 2I - (aq) PbI 2(s)

83 Note: if every ion in a complete ionic equation is a spectator ion, then no reaction occurs.

84 The net ionic equation includes only the ions and molecules directly involved in the reaction.

85 How can net ionic equations be used? They can be used to show similarities between large numbers of reactions A net ionic equation shows that more than one set of reactions can lead to the same net reaction.

86 What are the steps for writing net ionic equations?

87 1.) Write a balanced molecular equation for the reaction. 2.) Rewrite the equation to show the ions that form only strong electrolytes are written in ionic form. 3.) Identify and cancel spectator ions.

88 Sample Exercise 4.4 Write the net ionic equation for the precipitation reaction that occurs when solutions of calcium chloride and sodium carbonate are mixed.

89 Write the net ionic equation for the precipitation reaction that occurs when aqueous solutions of silver nitrate and potassium phosphate are mixed.

90 4.3 Acid Base Reactions

91 Acids and bases happen to be common electrolytes.

92 Acids

93 What are acids?

94 Acids are substances that ionize in aqueous solutions to form hydrogen ions, increasing the concentration of hydrogen ions in solution. Because hydrogen ions are just a proton, acids are known as proton donors.

95 The number of hydrogen ions produced depends on the number of hydrogen atoms present. Acids like HCl and HNO3 yield one hydrogen ion per molecule. Some acids such as H2SO4 yield 2 hydrogen ions.

96 What is the difference between a monoprotic an diprotic acid?

97 Acids that yield one hydrogen ion are monoprotic such as HCl HCl H + (aq) + Cl - (aq) Acids that yield two hydrogen ions are diprotic i.e., H 2 SO 4 H 2 SO 4 2H + (aq) + SO 4 2- (aq)

98 Bases

99 What are bases?

100 Bases are substances that accept H + ions thereby reducing the number of H + ions in solution. Bases produce OH - ions when dissolved in water. When bases are dissolved in water they release OH - and create more OH - ions by bonding to all of the available H + ions.

101 What are bases? Some bases do not contain OH - NH 3 is an example NH 3 + H 2 O NH OH -

102 Strong And Weak Acids and Bases

103 What are strong and bases?

104 Acids and bases that are strong electrolytes are called strong acids or bases.

105 What are weak acids and bases?

106 Acids and bases that are only partly ionized in solution.

107 Table 4.2 on page 122 lists the strong acids and bases. These should be committed to memory. What are the strong acids and bases.

108 ACIDS HClHClO 3 HNO 3 HBrHClO 4 H 2 SO 4 HI

109 Table 4.2 on page 122 lists the strong acids and bases. These should be committed to memory. What are the strong acids and bases. Bases The alkali group hydroxides Alkaline earth metals Ca, Sr and Ba hydroxides.

110 Identifying Strong and Weak Electrolytes

111 Is a soluble ionic compound a strong electrolyte, weak electrolyte or a nonelectrolyte?

112 All soluble ionic compounds are strong electrolytes.

113 How can you tell if a soluble molecular compound is a strong electrolyte, weak electrolyte or nonelectrolyte? All strong acids and bases are strong electrolytes All weak acids and bases are weak electrolytes All other soluble molecular compounds are nonelectrolytes.

114 The following diagrams represent aqueous solutions of three acids (HX, HY, and HZ) with water molecules omitted for clarity. Rank them from strongest to weakest.

115 Classify each of the following dissolved substances as a strong electrolyte, weak electrolyte, or nonelectrolyte: CaCl 2, HNO 3, C 2 H 5 OH (ethanol), HCOOH (formic acid), KOH.

116 Neutralization Reactions and Salts

117 What is a neutralization reaction?

118 When a solution of an acid and a base are mixed and the pH of the mixture is neither acidic or basic.

119 What are the products of a neutralization reaction?

120 The products of a neutralization reaction are always a salt and water. HCl + NaOH NaCl + H 2 O

121 What is the definition of a salt?

122 Any ionic compound whose cation comes from a base and whose anion comes from an acid.

123 What is the net ionic equation for all neutralization reactions?

124 H + (aq) + OH - (aq) H 2 O (l)

125 What type of reaction is a neutralization reaction? A double replacement (also known as a metathesis reaction or exchange reaction) Mg(OH) 2 + 2HCl MgCl 2 + 2H 2 O

126 Sample exercise 4.7 (a) Write a balanced molecular equation for the reaction between aqueous solutions of acetic acid (CH3COOH) and barium hydroxide, Ba(OH) 2. (b) Write the net ionic equation for this reaction.

127 (a) Write a balanced molecular equation for the reaction of carbonic acid (H 2 CO 3 ) and potassium hydroxide (KOH). (b) Write the net ionic equation for this reaction.

128 Acid Base Reactions with gas formation

129 There are two bases besides OH- that react with H+. Two of these include the sulfide and carbonate ions.

130 Both of these react with acids to form gases. 2HCl + Na 2 S H 2 S + 2NaCl HCl + NaHCO 3 NaCl + H 2 CO 3

131 Both NaHCO3 ( Sodium carbonate) and Na2HCO3 (Sodium Bicarbonate) are used as acid neutralizers and antacids. Read chemistry at work antacids.

132 Homework evens. Page 146

133 Section 4 Oxidation Reduction Reactions

134 What is corrosion? The conversion of a metal into a metal compound by a reaction between the metal and some substance in its environment.

135 What is corrosion? When a metal corrodes it lose electrons and forms cations Ca + 2HCl CaCl 2 + H I______________l Calcium is oxidized because it lost electrons

136 What is corrosion? Ca + 2HCl CaCl 2 + H l_______________l Hydrogen is reduced because it gained electrons.

137 What are oxidation – reduction (redox) reactions? Reactions in which electrons are transferred between reactants.

138 What is oxidation? When an atom becomes positively charged. When it has lost electrons Loss of electrons by a substance is called oxidation.

139 The term oxidation is used because the first reactions of this sort to be studied were reactions with oxygen.

140 What is reduction? Gain of electrons from a substance. The oxidation of one substance is always accompanied by the reduction of another substance.

141 Oxidation Numbers

142 What are oxidation numbers? The oxidation number of an atom in a substance is the actual charge of the atom if it is a monoatomic ion or it is the hypothetical charge assigned using a set of rules

143 Rules for assigning oxidation numbers 1.) For an atom in the elemental form, the oxidation number is always zero. H 2, Ca, O 2

144 Rules for assigning oxidation numbers 2.) For any monatomic ion, the oxidation number equals the charge on the ion K + = 1 + S 2- = 2 -

145 Rules for assigning oxidation numbers 3.) non-metals usually have negative oxidation numbers. a.) oxygen is usually 2 - w/ the exception of the peroxide ion (O 2 ) which has the oxidation number 1 -

146 Rules for assigning oxidation numbers b.) hydrogen has an oxidation number of 1 + when bonded to a nonmetal [(HCl) H 1 + ; Cl 1 - ] and has a oxidation of 1 - when bonded to a metal [ CaH 2 – Ca 2+, H 1 - ]

147 Rules for assigning oxidation numbers The sum of the oxidation numbers of all atoms in a nuetral compound is zero. The sum of the oxidation numbers in a polyatomic ion equals the charge of the ion.

148 Sample exercise 4.8 page 128 Determine the oxidation number of sulfur in each of the following: (a) H 2 S, (b) S 8, (c) SCl 2, (d) Na 2 SO 3, (e) SO 4 2–.

149 What is the oxidation state of the element in each of the following: (a) P 2 O 5 (b) NaH (c) Cr 2 O 7 2– (d) SnBr 4 (e) BaO 2

150 Oxidation of metals by acids and salts Some common types of redox reactions are combustion reactions and reactions between metals and acids or salts.

151 Oxidation of metals by acids and salts The common form of an acid reacting with a metal is A + BX AX + B Zn + 2HCl Zn Br 2 + H 2

152 What do we call these types of reactions and why are they classified as redox reactions? These reactions are called displacement or single replacement reactions. The ion in solution is displaced or replaced through the oxidation of an element.

153 Use the reaction between magnesium and hydrochloric acid to show that oxidation and reduction have occurred. Mg(s) + 2HCl MgCl 2 (aq) + H l___oxidized__l l_____reduced_____l

154 Write the net ionic equation for the reaction of magnesium and hydrochloric acid. Mg (s) + 2H + (aq) + 2Cl - Mg 2+ (aq) + 2Cl - + H 2(g) Cl - is a spectator ion. Mg (s) + 2H + (aq) Mg 2+ (aq) + H 2(g)

155 Metals can also be oxidized by aqueous solutions of various salts. Show the oxidation – reduction that occurs when iron reacts with nickel II nitrate

156 Fe (s) + Ni(NO 3 ) 3(aq) Fe(NO 3 ) 2(aq) + Ni (s) l______oxidized_____l l___________reduced______l NO 3 is the spectator ion. Net ionic equation Fe (s) + Ni 2+ (aq) Fe 2+ (aq) + Ni (s) l______oxidized_____l l_____reduced______l

157 Sample 4.9 Write the balanced molecular and net ionic equations for the reaction of aluminum with hydrobromic acid.

158 Write the balanced molecular and net ionic equations for the reaction between magnesium and cobalt(II) sulfate * What is oxidized and what is reduced in the reaction?

159 The Activity Series Different metals vary in the ease with which they are oxidized.

160 What is the activity series? The activity series is a list of metals arranged in order of decreasing ease of oxidation. The metals at the top of the table are the most easily oxidized. They react the most easily to form compounds.

161 What are active metals? The metals at the top of the activity series are the most easily oxidized metals.

162 Which are the noble metals? The metals are the bottom of the activity series. These metals are very stable and can be used to make coins and jewelry.

163 How can the activity series be used to predict the outcome of reactions? Any metal on the list can be oxidized by the ions of the an element below it. Cu + HCl No reaction Copper is not oxidized by hydrogen because hydrogen is not below copper Cu + AgNO 3 Ag + Cu(NO 3 ) 2 Copper is oxidized by silver because silver is below copper on the activity series.

164 Sample Exercise 4.10 Will an aqueous solution of iron(II) chloride oxidize magnesium metal? If so, write the balanced molecular and net ionic equations for the reaction.

165 Which of the following metals will be oxidized by Pb(NO 3 ) 2 : Zn, Cu, Fe?

166 Homework

167 Section 5 Concentrations of Solutions

168 Define concentration- The amount of solute dissolved in a given quantity of solvent.

169 What is Molarity? Molarity (M) expresses the concentration of a solution as the number of moles of solute in a liter of solution. Molarity (M) = moles of solute volume of soln.(L)

170 Sample 4.11 Calculate the molarity of a solution made by dissolving 23.4 g of sodium sulfate (Na 2 SO 4 ) in enough water to form 125 mL of solution.

171 Calculate the molarity of a solution made by dissolving 5.00 g of glucose (C 6 H 12 O 6 ) in sufficient water to form exactly 100 mL of solution.

172 Expressing the Concentration of an Electrolyte When an ionic compound dissolves, the relative concentrations of the ions introduced into the solution depend on the chemical formula.

173 Expressing the Concentration of an Electrolyte 1 mol of NaCl – 1 mole Na + ions 1 mole Cl - ions 1 mol of Na 2 SO mole Na + ions 1 mole SO 4 2- ions

174 Sample exercise 4.12 What are the molar concentrations of each of the ions present in a M aqueous solution of calcium nitrate?

175 What is the molar concentration of K+ ions in a M solution of potassium carbonate?

176 Interconverting Molarity, Moles and Volume Because molarity contains 3 quantities; molarity, moles and volume. Dimensional analysis can be used to find any of these values if we know the other two.

177 Calculate the number of moles of HNO 3 in 2.00L of a 0.200M HNO 3 # mol HNO 3 = ( 2 L HNO 3 ) (.200molHNO 3 ) 1 L HNO 3 =.4 mol HNO 3

178 Sample exercise 4.13 How many grams of Na 2 SO 4 are required to make L of M Na 2 SO 4 ?

179 How many grams of Na 2 SO 4 are there in 15 mL of 0.50 M Na 2 SO 4 ?

180 Dilutions

181 What are stock solutions? The concentrated solutions.

182 When solvent is added to dilute a stock solution the number of moles of solute before dilution is equal to the number of moles of solute after dilution.

183 To prepare 250mL of a M CuSO 4 from a stock of 1M CuSO 4 … 1 st determine the number of moles of CuSO 4 we will need in the dilute solution. (.250 L) (.10 mol/ 1L) =.0250 mol CuSO 4

184 To prepare 250mL of a M CuSO 4 from a stock of 1M CuSO 4 … Then determine the volume of stock solution needed L conc. Soln.=.025 mol CuSO 4 ( 1L/ 1 mole CuSO 4 ) =.025 L of 1 molar CuSO 4 = 25 mL Add 25 mL of 1 molar CuSO 4 to a 250 mL volumetric flask and bring up to volume.

185 To work the same problem quickly we can note Moles of solute in concentrated soln. = moles of solute in diluted soln. M conc. V conc = M dil. V dil ( 1M) ( V conc ) = (.1M) ( 250 mL) V conc = 25 mL

186 Sample exercise 4.14 How many milliliters of 3.0 M H 2 SO 4 are needed to make 450 mL of 0.10 M H 2 SO 4 ?

187 What volume of 2.50 M lead(II) nitrate solution contains mol of Pb 2+ ?

188 How many milliliters of 5.0 M K 2 Cr 2 O 7 solution must be diluted to prepare 250 mL of 0.10 M solution?

189 If 10.0 mL of a 10.0 M stock solution of NaOH is diluted to 250 mL, what is the concentration of the resulting stock solution?

190 Solution Stoichiometery and Chemical Analysis Recall that the coefficients in a balanced equation give the relative number of moles of reactants and products.

191 Sample exercise 4.15 How many grams of Ca(OH) 2 are needed to neutralize 25.0 mL of.10M HNO 3

192 How many grams of NaOH are needed to neutralize 20.0 mL of M H 2 SO 4 solution?

193 How many liters of M HCl(aq) are needed to react completely with mol of Pb(NO 3 ) 2 (aq), forming a precipitate of PbCl 2 (s)?

194 Titrations

195 What is a titration? A titration involves combining a sample of the solution with a reagent solution of known concentration called the standard solution.

196 What is the equivalence point? The point at which stoichiometrically equivalent quantities are brought together.

197 How does a chemist know when the equivalence point is reached? An indicator is used. The indicator will show pH changes when the color changes the acid has been nuetralized. The color change indicates the end point of the titration.

198 Sample exercise 4.16 The quantity of Cl– in a municipal water supply is determined by titrating the sample with Ag+. The reaction taking place during the titration is Ag+(aq) + Cl-(aq) AgCl(s) The end point in this type of titration is marked by a change in color of a special type of indicator. (a) How many grams of chloride ion are in a sample of the water if 20.2 mL of M Ag+ is needed to react with all the chloride in the sample? (b) If the sample has a mass of 10.0 g, what percent Cl– does it contain?

199 A sample of an iron ore is dissolved in acid, and the iron is converted to Fe 2+. The sample is then titrated with mL of M MnO4– solution. The oxidation-reduction reaction that occurs during titration is as follows: (a) How many moles of MnO4– were added to the solution? (b) How many moles of Fe2+ were in the sample? (c) How many grams of iron were in the sample? (d) If the sample had a mass of g, what is the percentage of iron in the sample?

200 One commercial method used to peel potatoes is to soak them in a solution of NaOH for a short time, remove them from the NaOH, and spray off the peel. The concentration of NaOH is normally in the range of 3 to 6 M. The NaOH is analyzed periodically. In one such analysis, 45.7 mL of M H 2 SO 4 is required to neutralize a 20.0-mL sample of NaOH solution. What is the concentration of the NaOH solution?

201 What is the molarity of an NaOH solution if 48.0 mL is needed to neutralize 35.0 mL of M H 2 SO 4 ?

202 (b) How many milliliters of 0.50 M Na 2 SO4 solution are needed to provide mol of this salt?


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