Chapter 04: Reactions in Aqueous Solution. Chemistry, 4 th Edition McMurry/Fay. Types of Chemical Reactions 01. Types of Chemical Reactions 02

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1 4 Chapter Reactions in Aqueous Solution Chemistry, 4 th Edition McMurry/Fay Dr. Paul Charlesworth Michigan Technological University Types of Chemical Reactions 01 Precipitation Reactions: A process in which an insoluble solid precipitate drops out of the solution. Most precipitation reactions occur when the anions and cations of two ionic compounds change partners. Pb(NO 3 ) 2 (aq) + 2 KI(aq) 2 KNO 3 (aq) + PbI 2 (s) Chapter 04 Slide 2 Types of Chemical Reactions 02 Acid Base Neutralization: A process in which an acid reacts with a base to yield water plus an ionic compound called a salt. The driving force of this reaction is the formation of the stable water molecule. HCl(aq) + NaOH(aq) NaCl(aq) + H 2 O(l) Chapter 04 Slide 3

2 Types of Chemical Reactions 03 Oxidation Reduction (Redox) Reaction: A process in which one or more electrons are transferred between reaction partners. The driving force of this reaction is the decrease in electrical potential. Mg(s) + I 2 (g) MgI 2 (s) Chapter 04 Slide 4 Types of Chemical Reactions 04 Metathesis Reactions: These are reactions where two reactants just exchange parts. AX + BY AY + BX HNO 3 (aq) + KOH(aq) KNO 3 (aq) + HOH(l) BaCl 2 (aq) + K 2 SO 4 (aq) BaSO 4 (s) + 2 KCl(aq) Chapter 04 Slide 5 Electrolytes in Solution 01 Why do ionic compounds conduct electricity when molecular ones generally do not? Chapter 04 Slide 6

3 Electrolytes in Solution 02 Electrolytes: Dissolve in water to produce ionic solutions. Nonelectrolytes: Do not form ions when they dissolve in water. Chapter 04 Slide 7 Electrolytes in Solution 03 Dissociation: The process by which a compound splits up to form ions in the solution. Chapter 04 Slide 8 Electrolytes in Solution 04 Strong Electrolyte: Total dissociation when dissolved in water. Weak Electrolyte: Partial dissociation when dissolved in water. Chapter 04 Slide 9

4 Electrolytes in Solution 05 Chapter 04 Slide 10 Solubility Rules & Precipitation 01 Allow you to predict whether a reactant or a product is a precipitate. Soluble compounds are those which dissolve to more than 0.01 M. There are three basic classes of salts: Chapter 04 Slide 11 Solubility Rules & Precipitation Salts which are always soluble: All alkali metal salts: Cs +, Rb +, K +, Na +, Li + All ammonium ion (NH 4+ ) salts All salts of the NO 3, ClO 3, ClO 4, C 2 H 3 O, and HCO 3 ions Chapter 04 Slide 12

5 Solubility Rules & Precipitation Salts which are soluble with exceptions: Cl, Br, I ion salts except with Ag +, Pb 2+, & Hg 2 2+ SO 4 ion salts except with Ag +, Pb 2+, Hg 2+ 2, Ca 2+, Sr 2+, & Ba 2+ Chapter 04 Slide 13 Solubility Rules & Precipitation Salts which are insoluble with exceptions: O & OH ion salts except with the alkali metal ions, and Ca 2+, Sr 2+, & Ba 2+ ions CO 3, PO 3 4, S, CrO 4, & SO 3 ion salts except with the alkali metal ions and the ammonium ion Chapter 04 Slide 14 Solubility Rules & Precipitation 05 Predict the solubility of: (a) CdCO 3 (b) MgO (c) Na 2 S (d) PbSO 4 (e) (NH 4 ) 3 PO 4 (f) HgCl 2 Predict whether a precipitate will form for: (a) NiCl 2 (aq) + (NH 4 ) 2 S(aq) (b) Na 2 CrO 4 (aq) + Pb(NO 3 ) 2 (aq) (c) AgClO 4 (aq) + CaBr 2 (aq) Chapter 04 Slide 15

6 Acid Base Concepts 01 Arrhenius Acid: A substance which dissociates to form hydrogen ions (H + ) in solution. Chapter 04 Slide 16 Acid Base Concepts 02 Arrhenius Base: A substance that dissociates in, or reacts with, water to form hydroxide ions (OH ). Chapter 04 Slide 17 Acid Base Concepts 03 Brønsted Acid: Can donate protons (H + ) to another substance in solution. Brønsted Base: Can accept protons (H + ) from another substance in solution. Chapter 04 Slide 18

7 Acid Base Concepts 04 Dissociation of Water: This equilibrium gives us the ion product of water. K w = K c = [H + ][OH ] = 1.0 x Chapter 04 Slide 19 Acid Base Concepts 05 Lewis Acid: Electron pair acceptor. e.g.: Al 3+, H +, BF 3. Lewis Base: Electron pair donor. e.g.: H 2 O, NH 3, O. Bond formed is called a coordinate bond. Chapter 04 Slide 20 Acid Base Concepts 06 Chapter 04 Slide 21

8 ph - A Measure of Acidity 01 The ph of a solution is defined as the negative logarithm of the hydrogen ion concentration (in mol/l). ph = log[h + ] ph + poh = 14 Acidic solutions: [H + ] > 1.0 x 10 7 M, ph < 7.00 Basic solutions: [H + ] < 1.0 x 10 7 M, ph > 7.00 Neutral solutions: [H + ] = 1.0 x 10 7 M, ph = 7.00 Chapter 04 Slide 22 ph - A Measure of Acidity 02 Calculate the ph of a HNO 3 solution having a hydrogen ion concentration of 0.76 M. The OH ion concentration of a blood sample is 2.5 x 10 7 M. What is the ph of the blood? Chapter 04 Slide 23 Neutralization Reactions Neutralization Reaction: produces salt & water. HA(aq) + MOH(aq) H 2 O(l) + MA(aq) Write ionic and net ionic equations for the following: (a) Ca(OH) 2 (aq) + 2 CH 3 CO 2 H(aq) (b) HBr(aq) + Ba(OH) 2 (aq) (c) HCl(aq) + NH 3 (aq) Chapter 04 Slide 24

9 Oxidation Reduction Reactions 01 Redox reactions are those involving the oxidation and reduction of species. Oxidation and reduction must occur together. They cannot exist alone. Chapter 04 Slide 25 Oxidation Reduction Reactions 02 Oxidation Is Loss (of electrons) Anode Oxidation Reducing Agent Chapter 04 Slide 26 Oxidation Reduction Reactions 03 Reduction Is Gain (of electrons) Cathode Reduction Oxidizing Agent Chapter 04 Slide 27

10 Oxidation Reduction Reactions 04 Assigning Oxidation Numbers: All atoms have an oxidation number regardless of whether it carries an ionic charge. 1. An atom in its elemental state has an oxidation number of zero. 2. An atom in a monatomic ion has an oxidation number identical to its charge. Chapter 04 Slide 28 Oxidation Reduction Reactions An atom in a polyatomic ion or in a molecular compound usually has the same oxidation number it would have if it were a monatomic ion. A. Hydrogen can be either +1 or 1. B. Oxygen usually has an oxidation number of 2. In peroxides, oxygen is 1. C. Halogens usually have an oxidation number of 1. When bonded to oxygen, chlorine, bromine, and iodine have positive oxidation numbers. Chapter 04 Slide 29 Oxidation Reduction Reactions The sum of the oxidation numbers must be zero for a neutral compound and must be equal to the net charge for a polyatomic ion. A. H 2 SO 4 2(+1) + (?) + 4( 2) = 0 net charge? = 0 2(+1) 4( 2) = +6 B. ClO 4 (?) + 4( 2) = 1 net charge? = 1 4( 2) = +7 Chapter 04 Slide 30

11 Oxidation Reduction Reactions Whenever one atom loses electrons (is oxidized), another atom must gain those electrons (be reduced). A substance which loses electrons (oxidized) is called a reducing agent. Its oxidation number increases. A substance which gains electrons (reduced) is called the oxidizing agent. Its oxidation number decreases. Chapter 04 Slide 31 Oxidation Reduction Reactions 07 Assign oxidation numbers to each atom in the following substances: A. CdS B. AlH 3 C. Na 2 Cr 2 O 7 D. SnCl 4 E. CrO 3 F. VOCl 3 G. V 2 O 3 H. HNO 3 I. FeSO 4 J. Fe 2 O 3 K. H 2 PO 4 L. MnO 4 M. Cr 2 O 7 Chapter 04 Slide 32 Oxidation Reduction Reactions 08 For each of the following, identify which species is the reducing agent and which is the oxidizing agent. Ca(s) + 2 H + (aq) Ca 2+ (aq) + H 2 (g) 2 Fe 2+ (aq) + Cl 2 (aq) 2 Fe 3+ (aq) + 2 Cl (aq) SnO 2 (s) + 2 C(s) Sn(s) + 2 CO(g) Sn 2+ (aq) + 2 Fe 3+ (aq) Sn 4+ (aq) + 2 Fe 2+ (aq) Chapter 04 Slide 33

12 Activity Series of Elements 01 Chapter 04 Slide 34 Activity Series of Elements 02 Activity series looks at the relative reactivity of a free metal with an aqueous cation. Fe(s) + Cu 2+ (aq) Fe 2+ (aq) + Cu(s) Zn(s) + Cu 2+ (aq) Zn 2+ (aq) + Cu(s) Cu(s) + 2 Ag + (aq) 2 Ag(s) + Cu 2+ (aq) Mg(s) + 2 H + (aq) Mg 2+ (aq) + H 2 (g) Chapter 04 Slide 35 Activity Series of Elements 03 Given the following three reactions, determine the activity series for Cu, Zn, & Fe. Fe(s) + Cu 2+ (aq) Fe 2+ (aq) + Cu(s) Zn(s) + Cu 2+ (aq) Zn 2+ (aq) + Cu(s) Fe(s) + Zn 2+ (aq) NR Chapter 04 Slide 36

13 Metathesis & Net Ionic Reactions 01 Three basic equations are used for reactions: Molecular Equation: All reactants and products are written in molecular form. Ionic Equation: All dissolved strong electrolytes are written as the dissociated ions. Net Ionic Equation: All ions that are identical on both sides are deleted. Chapter 04 Slide 37 Metathesis & Net Ionic Reactions 02 Molecular Equation: All reactants and products are written in molecular (non-dissociated) form along with their phases. Pb(NO 3 ) 2 (aq) + 2 KCl(aq) PbCl 2 (s) + 2 KNO 3 (aq) 2 HCl(aq) + Cu(OH) 2 (s) CuCl 2 (aq) + 2 HOH(l) C 2 H 3 O 2 H(aq) + KOH(aq) KC 2 H 3 O 2 (aq) + HOH(l) Chapter 04 Slide 38 Metathesis & Net Ionic Reactions 03 Ionic Equation: All dissolved strong electrolytes in the molecular equation are broken into their ions. Pb NO K Cl PbCl 2 (s) + 2 K NO 3 2 H Cl + Cu(OH) 2 (s) Cu Cl + 2 HOH(l) C 2 H 3 O 2 H(aq) + K + + OH K + + C 2 H 3 O + HOH(l) Chapter 04 Slide 39

14 Metathesis & Net Ionic Reactions 04 Net Ionic Equation: Spectator ions that occur on both sides are cancelled to give only those species which undergo change. Pb Cl PbCl 2 (s) 2 H + + Cu(OH) 2 (s) Cu HOH(l) C 2 H 3 O 2 H(aq) + OH C 2 H 3 O + HOH(l) Chapter 04 Slide 40 Metathesis & Net Ionic Reactions 05 Write net ionic equations for the following reactions: 2 AgNO 3 (aq) + Na 2 CrO 4 (aq) Ag 2 CrO 4 (s) + 2 NaNO 3 (aq) H 2 SO 4 (aq) + MgCO 3 (s) H 2 O(l) + CO 2 (g) + MgSO 4 (aq) Chapter 04 Slide 41 Balancing Redox Reactions 01 Half-Reaction Method: Allows you to focus on the transfer of electrons. This is important when considering batteries and other aspects of electrochemistry. The key to this method is to realize that the overall reaction can be broken into two parts, or halfreactions. Chapter 04 Slide 42

15 Balancing Redox Reactions 02 Balance for an acidic solution: MnO 4 (aq) + Br (aq) Mn 2+ (aq) + Br 2 (aq) The steps involved follow the same basic procedure as described for the oxidation-number method. 1. Determine oxidation and reduction halfreactions: Oxidation half-reaction: Br (aq) Br 2 (aq) Reduction half-reaction: MnO 4 (aq) Mn 2+ (aq) Chapter 04 Slide 43 Balancing Redox Reactions Balance for atoms other than H and O: Oxidation: 2 Br (aq) Br 2 (aq) Reduction: MnO 4 (aq) Mn 2+ (aq) 3. Balance for oxygen by adding H 2 O: Oxidation: 2 Br (aq) Br 2 (aq) Reduction: MnO 4 (aq) Mn 2+ (aq) + 4 H 2 O(l) Chapter 04 Slide 44 Balancing Redox Reactions Balance for hydrogen by adding H + : Oxidation: 2 Br (aq) Br 2 (aq) Reduction: MnO 4 (aq) + 8 H + (aq) Mn 2+ (aq) + 4 H 2 O(l) 5. Balance for charge by adding electrons (e ): Oxidation: 2 Br (aq) Br 2 (aq) + 2 e Reduction: MnO 4 (aq) + 8 H + (aq) + 5 e Mn 2+ (aq) + 4 H 2 O(l) Chapter 04 Slide 45

16 Balancing Redox Reactions Balance for numbers of electrons by multiplying: Oxidation: 5[2 Br (aq) Br 2 (aq) + 2 e ] Reduction: 2[MnO 4 (aq) + 8 H + (aq) + 5 e Mn 2+ (aq) + 4 H 2 O(l)] 7. Combine and cancel to form one equation: Oxidation: 10 Br (aq) 5 Br 2 (aq) + 10 e Reduction: 2 MnO 4 (aq) + 16 H + (aq) + 10 e 2 Mn 2+ (aq) + 8 H 2 O(l) 2 MnO 4 (aq) + 10 Br (aq) + 16 H + (aq) 2 Mn 2+ (aq) + 5 Br 2 (aq) + 8 H 2 O(l) Chapter 04 Slide 46 Balancing Redox Reactions 06 Balancing for Basic Solution: MnO 4 (aq) + SO 3 (aq) MnO 4 (aq) + SO 4 (aq) 7. Balance for acidic solution. 2 MnO 4 (aq) + SO 3 (aq) + H 2 O(l) 2 MnO 4 (aq) + SO 4 (aq) + 2 H + (aq) Chapter 04 Slide 47 Balancing Redox Reactions Add OH to neutralize the H + ions from the acidic balancing. 2 MnO 4 (aq) + SO 3 (aq) + H 2 O(l) + 2 OH (aq) 2 MnO 4 (aq) + SO 4 (aq) + 2 H + (aq) + 2 OH (aq) 9. Reduce to simplest form. 2 MnO 4 (aq) + SO 3 (aq) + 2 OH (aq) 2 MnO 4 (aq) + SO 4 (aq) + H 2 O(l) Chapter 04 Slide 48

17 Balancing Redox Reactions 13 Balance the following for acidic and basic solution: ClO (aq) + Cr(OH) 4 (aq) CrO 4 (aq) + Cl (aq) NO 3 (aq) + Cu(s) NO(g) + Cu 2+ (aq) Fe(OH) 2 (s) + O 2 (g) Fe(OH) 3 (s) MnO 4 (aq) + IO 3 (aq) MnO 2 (s) + IO 4 (aq) Cr 2 O 7 (aq) + Fe 2+ (aq) CrO 4 (aq) + Cl (aq) Chapter 04 Slide 49

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